QD 

45- 


SB    flOM 


A  Laboratory  Manual 
of  General  Chemistry 


A  LABORATORY  MANUAL 


OF 


GENERAL    CHEMISTRY 


BY 


W.  S.  HENDRIXSON, 

Professor  of  Chemistry  in  Grinnell  College. 


GRINNELL,  IOWA, 

1918 


ALL   RIGHTS   RESERVED 


PREFACE. 

This  the  fourth  edition  of  the  author's  Experiments  in  General 
Chemistry  is  printed,  as  its  predecessors  have  been,  primarily  for  the 
use  of  students  in  general  chemistry  in  Grinnell  College.  This  fact 
may  explain  certain  departures  from  custom  in  the  preparation  of 
such  books,  such  as  suggestions  to  teachers  and  detailed  descriptions 
of  apparatus  and  its  manipulation.  As  a  matter  of  fact  apparatus  at  all 
complicated  is  not  only  shown  by  cuts,  but  it  is  set  up  on  the  lecture 
table  and  many  experiments  for  any  period  are  there  carried  through 
before  the  students  enter  the  laboratory.  Some  apparatus  is  even  set 
up  in  the  laboratory  and  left  there  for  inspection  during  the  laboratory 
perio'd. 

Not  satisfied  to  use  the  same  set  of  even  his  own  experiments  year 
after  year  and  wishing  to  provide  new  laboratory  work  for  classes  of 
students  who  have  taken  chemistry  in  the  high  school,  the  writer  has 
provided  for  more  laboratory  work  than  can  be  done  in  a  three-  or 
four:hour  course  of  one  year. 

In  this  book  an  attempt  is  made  to  connect  rationally  general 
chemistry  and  qualitative  analysis.  Students  who  complete  a  first 
year  course  in  chemistry  should  have  some  knowledge  of  qualitative 
analysis,  but  it  should  not  be  permitted  to  take  the  place  of  general 
chemistry  in  the  second  half  year,  which  is  usually  devoted  to  study  of 
the  metals.  Qualitative  analysis  ought  to  be  a  development  from  the 
general  chemistry  to  which  it  gives  point,  and  its  introduction  as  an 
outgrowth  of  the  general  chemistry  greatly  stimulates  the  student's 
interest  in  both  subjects.  In  this  book  tests  for  acids  and  other  com- 
pounds are  given  in  the  study  of  the  non-metals,  and  a  system  for  the 
detection  of  acids  is  given  after  the  study  of  the  non-metals  has  been 
completed.  In  the  study  of  the  metals  emphasis  is  placed  on  properties 
that  are  of  analytical  significance,  though  other  facts  are  not  neg- 
lected. After  each  group  of  metals  has  been  studied  their  separation 
is  taken  up,  and  the  work  is  extended  as  rapidly  as  the  student's  ex- 
perience justifies  it,  to  the  detection  of  both  metals  and  acid  radicals 
in  "unknowns."  The  scheme  of  qualitative  analysis  as  outlined  is  not 
supposed  to  be  complete  but  is  meant  to  serve  as  an  introduction  to 
the  subject  and  a  preparation  to  the  more  rigorous  course  in  quali- 
tative analysis  the  following  year. 


387354 


TO  THE  STUDENT. 

I.— On  coming  to  the  laboratory  the  first  day  find  the  number  of 
your  desk,  get  the  key  or  combination  and  a  printed  list  of  the  appa- 
ratus the  desk  should  contain.  Verify  the  apparatus,  asking  the  names 
of  the  things  you  do  not  know.  Make  sure  that  apparatus  is  wanting 
before  calling  for  it.  Present  broken  or  faulty  apparatus  for  exchange. 

II. — After  using  apparatus,  clean  it  and  return  it  to  the  desk. 
Clean  the  top  of  the  desk  and  lock  it  before  leaving.  Have  old  cloths 
for  cleaning  and  a  towel  for  the  hands.  In  grading  notes  you  will  be 
held  responsible  for  the  condition  of  the  desk  without  and  within  anfl 
for  the  bottles  on  its  shelves,  which  should  always  be  kept  in  the 
same  order. 

III. — Provide  an  approved  note  book  of  the  sort  shown  in  the  lec- 
ture room.  Leave  the  first  leaf  blank  and  enough  margins  to  permit 
corrections  by  the  instructor. 

Begin  every  experiment  on  a  new  page.  Write  notes  in  the  labor- 
atory in  full  and  do  not  copy  them.  Only  original  notes  made  while 
doing  the  laboratory  work  are  of  value.  The  notes  on  each  experiment 
should  have  a  suitable  heading  and  should  have  the  same  number  as 
the  experiment.  Paragraph  to  suit  the  subject  matter.  Present  the 
note  book  for  grading  and  criticism  before  leaving  the  laboratory. 

IV.— Ascertain  in  advance  what  the  laboratory  work  is  to  be  for 
each  period  and  prepare  for  it  before  coming  to  the  laboratory  by  use 
of  the  text  or  reference  book.  The  subject  will  be  announced  in  the 
class  room,  and  references  to  books  will  be  given.  Bring  the  text- 
book to  the  laboratory. 

V. — Every  experiment  is  intended  to  teach  something  of  value. 
Try  to  find  out  what  it  is  by  yourself,  using  observation  and  the  text- 
book, then  ask  the  instructor.  State  results  in  the  notes.  Do  not  con- 
fine these  efforts  merely  to  answering  the  questions  of  the  laboratory 
book.  For  example  after  equations  have  been  studied  in  class  all 
equations  for  reactions  in  the  laboratory  work  should  be  written 
whether  asked  for  or  not.  Write  them  where  they  belong  scattered 
through  the  notes  and  do  not  write  them  in  mass  at  the  end  of  the 
notes. 

VI. — When  using  chemicals  replace  the  covers  or  stoppers  of  the 
containing  vessels.  Do  not  throw  stoppers  upon  the  desk.  Return  all 
bottles  to  their  places  on  the  shelves.  Replace  all  weights  in  the 
boxes. 

VII. — Do  not  wash  pieces  of  apparatus  with  distilled  water  but. 
with  hydrant  water.  If  they  are  well  drained  it  will  not  be  necessary 
in  most  cases  even  to  rinse  them  with  distilled  water.  On  the  other 


LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY  3 

hand  always  have  distilled  water  in  your  wash  bottle  and  use  it  in 
making  all  solutions. 

VIII. — Throw  no  solid  matter  into  the  sinks  but  into  the  jars  pro- 
vided for  that  purpose.  On  the  other  hand  wash  apparatus  at  the 
sinks  and  do  not  throw  large  amounts  of  water  into  the  jars. 

DIRECTIONS  FOE  LABORATORY  WORK. 
GENERAL  MANIPULATION. 

1. — The  Bunsen  Burner:  Take  apart  the  burner  and  study  its  con- 
struction. Determine  how  to  regulate  the  supplies  of  gas  and  air.  Put 
it  together  and  turn  on  the  gas  at  the  cock  and  regulate  the  supply  by 
the  screw  on  the  burner.  Regulate  the  gas  and  air  so  as  to  secure  a 
non-luminous  flame.  Too  much  air  may  cause  the  flame  to  blow  out, 
"snap  back"  or  burn  with  much  noise.  A  long  green  hissing  flame  in- 
dicates that  the  gas  is  burning  also  at  the  base.  Turn  out  the  flame, 
reduce  the  amount  of  air,  light  again.  A  flame  about  3  inches  long  is 
usually  sufficient. 

The  ordinary  flame  is  used  to  heat  test  tubes  directly,  or  to  heat 
such  as  flasks  and  beakers  placed  upon  wire  gauze,  of  which  nichrome 
gauze  is  best.  The  crown  top  is  used  to  heat  beakers  and  flasks  with- 
out protection  of  gauze. 

The  Wing  Top :  Put  the  wing  top  on  the  burner,  turn  on  the  gas 
and  regulate  gas  and  air  so  as  to  secure  a  just  non-luminous  flame 
about  as  high  as  broad.  Too  much  air  is  objectionable.  Such  a  flame 
should  be  used  exclusively  for  bending  tubing. 

'2.— Breaking  and  Bending  Glass  Tubing:  Draw  once  a  sharp,  three- 
cornered  file  across  the  piece  of  tubing  where  you  wish  to  break  it. 
Hold  the  tube  with  both  hands,  with  thumbs  together  and  opposite  the 
scratch.  A  slight  pull  will  break  the  tube  squarely  at  the  scratch. 

To  bend  glass  tubing  always  use  the  wing  top.  No  one  but  an  ex- 
pert can  make  a  gcod  bend  with  the  flame  of  the  ordinary  burner.  A 
bend  so  made  will  be  uneven,  crinkled  and  the  tube  is  likely  to  break 
at  that  point. 

Hold  the  tube  above  the  middle  of  the  wing  flame,  having  its 
length  in  direction  of  the  breadth  of  the  flame.  Revolve  the  tube  and 
move  it  back  and  forth  in  the  direction  of  its  length  until  it  softens 
and  begins  to  yield  to  a  slight  pressure,  then  bend  it  slowly  as  de- 
sired. If  it  shows  a  tendency  to  collapse  or  flatten  at  any  point,  stop 
bending  at  that  point,  heat  at  one  side  or  the  other  and  there  complete 
the  bend.  The  tendency  to  collapse  shows  that  the  bend  is  being  made 
too  short.  Practice  with  scraps  of  tubing  until  you  can  make  a  good 
bend,  and  then  make  the  tubes  like  the  samples  shown  in  fig.  1. 
They  represent  the  tubes  most  used  in  this  course,  and  should  be  kept 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


Fig.  1 

throughout  the  year  and  not  cut  up  to  make  other  things.  Heat  the 
ends  of  the  tubes  to  redness  in  the  ordinary  flame  to  fuse  down  the 
sharp  edges.  Other  glass  working  will  be  shown  in  the  class  room  as 
occasions  arise. 

If  there  is  no  wash  bottle  in  the  desk  construct  one  as  shown  in 
fig.  1.  This  wash  bottle  should  always  contain  a  supply  of  distilled 
water  and  should  not  be  torn  down  to  use  the  tubes  or  flask  for  any 
other  purpose. 

3. — Weighing.  The  following  may  be  applied  to  all  balances  and 
weighings,  but  will  be  supplemented  by  the  instructors  when  weigh- 
ings of  great  accuracy  are  required. 

Never  place  chemicals,  other  than  pieces  of  metals,  directly  upon 
the  pans  of  the  balance,  but  in  some  suitable  containing  vessel.  When 
only  moderate  accuracy  is  required,  balanced  papers  may  be  used,  but 
not  where  accurate  weighings  are  called  for.  In  the  latter  case  two 
methods  are  good.  Place  a  dish  on  the  left  hand  pan  and  weigh  it  ac- 
curately. Place  in  it  about  the  desired  amount  of  substance  and  weigh 
again.  Of  course  the  difference  between  the  weights  gives  the  weight 
of  the  substance.  The  scales  rarely  "balance"  or  the  pointer  rarely 
stands  at  the  zero  point.  This  error,  however,  is  eliminated  by  the 
above  method.  Instead  of  the  dish,  the  substance  may  be  weighed  in 
a  corked  tube,  then  the  desired  amount  may  be  taken  out  and  the  tube 
weighed  again.  Of  course  in  both  weighings  the  tube  must  be  in  the 
same  pan,  always  the  left. 

When  equilibrium  has  been  attained,  that  is,  when  the  pointer 
makes  excursions  equidistant  right  and  left  from  the  center,  count  the 
weights  without  removing  them  from  the  pan,  record  the  weight  in  the 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  5 

note  bcok  and  verify  by  counting  the  weights  again  as  they  are  re- 
moved to  the  box.  Write  the  weight  in  the  book  as  one  number.  If 
the  weights  are  5  g.,  2  g.,  200  mg.,  50  mg.,  write  the  total  at  once  7.25 
grains.  Unless  the  nature  of  the  experiment  requires  it,  it  is  a  mis- 
take to  try  to  weigh  a  definite  amount  of  a  substance.  For  example,  if 
the  directions  say  "Weigh  accurately  about  5  grams  of  the  substance," 
do  not  try  to  get  just  five  grams,  but  take  about  five  grams  and  weigh 
it  accurately. 

4. — Pleasuring  Volumes  of  Liquids:  Study  a  pipet  and  a  gradu- 
ated cylinder.  Why  is  the  cylinder  graduated  from  the  top  downward 
and  from  the  bottom  upward? 

Weigh  accurately  a  small  flask  on  the  horn-pan  scales  observing 
the  directions  given  under  weighing.  Now  put  a  small  piece  of  rub- 
ber tubing  on  the  larger  end  of  the  pipet,  by  suction  fill  it  with  dis- 
tilled water  to  above  the  mark  and  pinch  the  rubber  tube.  Carefully 
let  out  the  water  to  the  mark  and  then  run  the  remainder  into  the 
flask  and  weigh  again.  Find  the  weight  of  the  water,  the  weight  of 
1  cc.  and  the  error  in  the  pipet  according  to  your  work.  Why  in  this 
experiment  does  it  make  little  difference  whether  the  balance  with 
pans  empty  was  in  exact  equilibrium  or  not? 

OXYGEX. 

5. — Preparation  of  Oxygen:  (Study  text-book  in  advance).  Many 
substances  containing  it  give  off  all  or  a  part  of  their  oxygen  when 
heated.  The  following  are  illustrations: 

In  a  small  test  tube  heat  about  half  a  gram,  estimated,  of  potas- 
sium chlorate.  It  melts  and  then  seems  to  boil  owing  to  the  evolution 
of  oxygen.  Light  a  splinter  or  tooth  pick,  blow  out  flame,  lower  glow- 
ing end  into  the  test  tube.  It  should  rekindle. 

In  another  small  tube,  preferably  one  made  of  "hard  glass"  tubing 
heat  a  little  mercury  oxide  persistently  and  test  with  glowing  splinter. 
Note  sublimed  mercury  on  wall  of  tube.  By  persistent  heating  all  the 
mercury  oxide  may  be  decomposed  into  mercury  and  oxygen. 

As  in  the  case  of  mercury  oxide  heat  a  little  manganese  dioxide 
and  test  for  oxygen.  Only  one-third  of  the  oxygen  is  given  off. 

Oxygen  may  be  obtained  also  by  heating  lead  dioxide,  potassium 
permanganate,  potassium  perchlorate,  barium  dioxide  to  redness. 

•  All  these  instances  illustrate  chemical  change,  and  the  nature  of 
the  change  is  that  of  decomposition. 

({.—  Preparation  of  Oxygen,  Laboratory  Method:  (a)  Read  through 
the  experiment  and  have  everything  required  within  reach.  Set  up 
the  apparatus  as  shown  in  fig.  2,  having  trough  filled  with  water  to 
about  2  inches  above  the  shelf.  Place  the  delivery  tube  well  to  the 
farther  side  of  the  trough  so  as  not  to  be  obliged  to  reach  over  it. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


Fig.  2. 

Fill  two  jars*  with  water  by  immersing  in  the  trough  with  mouth 
slightly  upward,  inverting  them  and  placing  upon  the  shelf  as  shown. 
Direct  the  base  of  the  retort  stand  backward  so  that  it  will  not  be  in 
the  way  of  the  burner.  (See  fig.  2).  There  it  no  need  of  clamping  the 
test  tube  tightly  so  as  to  endanger  breaking  it.  Have  the  clamp  near 
its  mouth.  Why? 

Weigh  approximately  on  platform  scales  8  grams  potassium  chlor- 
ate and  4  grams  powdered  manganese  dioxide,  mix  them  on  paper  with 
a  spatula  and  from  the  paper  slide  the  mixture  into  the  test-tube. 

Now  heat  the  substance  in  the  tube  slowly  and  evenly,  best  by 
holding  the  burner  in  the  hand  and  moving  it  back  and  forth  along  the 
tube,  so  as  to  avoid  over-heating  and  softening  the  tube  at  any  one 
point.  So  regulate  the  heat  that  the  gas  shall  be  coming  off  slowly 
when  the  jar  is  about  full.  When  the  jar  is  full  slide  it  to  one  side 
and  place  another  jar  over  the  mouth  of  the  tube,  then  place  the  cover 
securely  on  the  full  jar  keeping  its  mouth  under  water,  remove  it  *o 
the  desk  and  clamp  on  the  cover.  Fill  five  jars  or  bottles.  When  the 
gas  ceases  to  come  off  remove  the  stopper  of  the  tube  to  prevent  a 
back  flow  of  the  water.  Allow  the  tube  to  cool  and  meantime  use  the 
oxygen  in  7.  Return  then  to  (b). 

(b)  Prepare  a  funnel  with  filter  paper.  Fold  a  filter  in  halves, 
then  in  fourths,  open  out  one  thickness  making  a  cone,  press  down 
evenly  into  a  funnel  and  wet  it  with  water  to  make  it  hold  its  shape. 
Fill  half  full  of  distilled  water  the  tube  used  to  heat  the  potassium 
chlorate  and  manganese  dioxide,  and  when  the  residue  is  wet  through- 
out by  shaking  heat  to  boiling  and  filter  by  pouring  the  contents  into 
the  filter,  being  careful  rot  to  run  it  over,  and  catch  the  clear  liquid, 
filtrate,  in  a  large  test  tube  or  beaker.  If  not  clear  run  it  through 
again.  To  a  little  of  the  clear  filtrate  in  a  clean  test  tube  add  a  few 

*  Rubber  seal,  pint  fruit  jar. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  7 

drops  of  solution  of  silver  nitrate.  Now  add  a  few  drops  of  dilute  ni- 
tric acid.  The  white  precipitate  is  silver  chloride.  The  formation  of 
this  white  precipitate  by  silver  nitrate  in  the  presence  of  nitric  acid  is 
a  much  used  test  for  a  chloride.  Make  solutions  of  potassium  chloride 
and  potassium  chlorate  from  the  shelf  bottles,  add  a  little  dilute  ni- 
tric acid  to  each,  then  silver  nitrate.  Which  gives  silver  chloride? 
What  does  your  clear  filtrate  seem  to  contain?  Evaporate  a  little  of 
it  to  dryness  in  a  porcelain  dish.  Compare  the  appearance  and  taste 
of  residue  with  potassium  chloride  and  chlorate  from  shelf.  What 
chemical  change  took  place  on  heating  potassium  chlorate?  Was  the 
manganese  dioxide  changed  chemically  on  heating?  For  the  use  of 
manganese  dioxide  as  a  "catalyzer"  see  (8). 

7. — Properties  of  Oxygen:  Place  about  half  a  gram  of  sulfur  in 
a  deflagration  spoon,  best  covered  with  a  bit  of  asbestos  paper,  ignite 
the  sulfur  and  lower  it  into  a  jar  of  oxygen.  When  the  combustion 
ceases,  add  a  little  distilled  water  to  the  jar  and  shake.  Test  the  water 
with  blue  litmus  paper.  The  change  to  red  shows  acid  formed  by  the 
union  of  water  with  sulfur  dioxide  produced  by  the  combustion. 

Ignite  a  piece  of  charcoal  on  a  clean  spoon  and  lower  into  a  jar 
of  oxygen  containing  a  little  water.  Cover  jar,  shake  and  test  the 
water  with  blue  litmus  paper.  A  faint  red  should  be  given  the  paper 
showing  the  formation  of  carbonic  acid.  Add  to  the  jar  clear  lime 
water*  and  shake.  The  milky  appearance  is  due  to  calcium  carbon- 
ate, produced  by  calcium  hydroxide  in  the  lime  water  and  carbon  di- 
oxide. 

Place  a  little  red  phosphorus  on  asbestos  paper  in  the  spoon,  light 
it  and  quickly  lower  into  oxygen.  Test  the  water  in  the  jar  with  blue 
litmus  paper.  Phosphorus  pentoxide  is 'first  formed  and  this  unites 
with  the  water  forming  phosphoric  acid.  What  is  an  element?  All 
three  of  the  substances  burned  are  elements.  These  are  non-metals, 
and  such  on  combustion  usually  form  oxides  that  are -acidic  or  unite 
with  water  to  form  acids. 

Heat  the  end  of  a  piece  of  picture  wire,  dip  into  powdered  sulfur 
for  an  instant.  A  portion  of  the  S  will  adhere  and  burn.  Lower  it 
into  a  jar  of  O.  The  sulfur  burns  and  then  the  iron  burns  with  a 
shower  of  sparks. 

Without  removing  the  kerosene  which  adheres  to  it  place  a  bit 
of  sodium  as  large  as  a  grain  of  wheat  in  a  spoon,  ignite  the  kerosene 
and  lower  into  a  jar  of  oxygen.  The  kerosene  should  burn  and  ignite 

*If  lime  water  is  not  present  in  the  laboratory  in  quantity,  it  may 
easily  be  prepared  by  shaking  a  little  slaked  lime  in  a  jar  of  water 
for  some  time,  letting  the  lime  settle  and  filtering.  The  funnel  should 
be  placed  in  a  flask  to  protect  the  filtrate  from  carbon  dioxide  in  the 
air. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

the  sodium.  Shake  the  bottle  and  test  the  water  with  red  litmus  pa- 
per. The  so-called  alkali  metals  such  as  sodium  burn  and  form  oxides 
which  unite  with  water  forming  an  alkali. 

The  burning  of  these  elements  forming  compounds,  oxides,  illus- 
trates chemical  change,  and  chemical  reactions  of  the  type  called  di- 
rect combination  or  synthesis. 

8.  Manganese  Dioxide  as  a  Catalyzer:  In  a  test  tube  heat  about 
1  gram  of  potassium  chlorate  till  it  melts  and  begins  to  give  off  oxy- 
gen. Without  cooling  drop  a  very  little  manganese  dioxide  upon  the 
fused  salt.  Is  there  much  change  in  the  rate  of  giving  off  oxygen? 
Compare  temperature  required  when  oxygen  is  prepared  from  a  mix- 
ture of  potassium  chlorate  and  manganese  dioxide,  and  from  potassium 
chlorate  alone.  The  next  experiment  illustrates  further  the  influence 
of  catalyzers,  and  also  a  very  good  method  of  preparing  oxygen. 

9. — Oxygen  from  Sodium  Dioxide  (Peroxide) :  In  each  of  three  dry 
test  tubes  place  about  a  gram  of  sodium  dioxide.  In  one  place  also  a 
very  little  fine  copper  oxide  and  in  another  powdered  manganese  di- 
oxide. In  each  of  the  three  pour  about  5c.c.  of  water  and  note  rates  at 
which  oxygen  is  given  off. 

To  prepare  a  larger  volume  of  oxygen  by  this  method  set  up  the 
apparatus  as  in  fig.  10,  but  use  a  dry  flask  instead  of  the  test  tuba. 
Place  in  flask  about  10  grams  of  sodium  dioxide,  then  about  half  a 
gram  of  powdered  copper  oxide  or  manganese  dioxide,  shake  to  mix 
and  spread  the  mixture  evenly  over  the  bottom  of  the  flask.  Fill  the 
funnel  with  water  and  by  means  of  the  pinch  cock  let  it  run  in  slowly 
so  as  to  maintain  a  suitable  flow  of  oxygen.  No  heating  is  necessary. 
The  oxygen  may  be  tested  with  a  glowing  splinter.  Sodium  hydroxide, 
NaOH,  is  formed  at  the  same  time.  Test  a  little  of  the  solution  in 
the  flask  with  red  litmus  or  turmeric  paper.  Test  sodium  hydroxide 
from  the  shelf  with  the  same  paper. 

HYDROGEX. 

11. —  (a).  In  as  many  test  tubes  place  bits  of  magnesium  turnings, 
aluminium  turnings  or  wire,  zinc,  iron  filings,  or  other  finely  divided 
iron  such  as  steel  wool;  tin,  copper,  lead.  To  each  tube  add  dilute 
hydrochloric  acid.  In  what  cases  is  hydrogen,  a  colorless  gas,  given 
off?  In  cases  where  it  is  not,  add  a  few  drops  of  concentrated  hydro- 
chloric acid.  If  some  metals  still  resist  warm  the  acid. 

(b).  In  clean  tubes  and  fresh  metals  try  the  same  series  with  di- 
lute sulfuric  acid  warming  where  necessary.  In  text-book  refer  to 
"electro-motive  series."  From  the  more  electro-positive  end  of  the 
series  the  metals  give  off  H  with  decreasing  readiness  till  hydrogen 
itself  is  reached.  The  metals  beyond  or  below  H  do  not  give  H  with 
acids.  That  is,  the  metals  that  are  more  electro-positive  than  H  dis- 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


place  it  from  solution.     There  are  secondary  reasons  why  some  met- 
als above  H  do  not  evidently  replace  hydrogen.     Lead  is  an  example, 
(c).    For  the  action  cf  nitric  acid  on  metals  see  52. 
12.— Preparation  of  Hydrogen.     Caution:     Mixtures   of  hydrogen 
and  oxygen  and  hydrogen  and  air  are  dangerously  explosive.     In  the 

preparation  of  hydrogen  in 
this  experiment  and  else- 
where have  no  flame  near  the 
delivery  tube  until  all  air  is 
expelled  from  the  flask. .  The 
gas  should  be  lighted  the  first 
time  with  a  test  tube  of  burn- 
ing gas.  Should  the  stopper 
of  the  flask  be  removed  so  as 
to  permit  the  entrance  of  air. 
Fig.  3.  the  same  precaution  must  be 

used  in  lighting  the  gas  the  second  time. 

In  a  small  flask  fitted  as  in  Fig.  3  place  about  25  grams  of  granu- 
lated zinc  and  about  50c.c.  of  dilute  sulphuric  acid.  Add  a  few  drops 
of  a  solution  of  copper  sulphate  if  the  action  proceeds  too  slowly  when 
sulfuric  acid  is  used.  To  get  the  crystals  mentioned  below  it  is  neces- 
sary to  u~e  rrulphuric  acid.  In  performing  the  experiment  subsequent- 
ly, inrtead  of  dilute  sulphuric  acid  the  zinc  may  be  covered  with  water 
and  then  strong  hydrochloric  acid  cautiously  added  through  the  fun- 
nel tube  until  the  gas  comes  off  freely.  Collect  the  gas  in  jars  reject- 
ing the  first  jar  full.  Why?  Collect  two  jars  and  then  light  the  gas 
with  a  test  tube  of  burning  gas.  That  is,  hold  the  delivery  tube  up- 
ward, and  place  over  its  end  a  small  test  or  specimen  tube.  The  hy- 
drogen will  rise  and  force  the  air  downward.  When  the  tube  has  had 
time  to  fill  move  it  mouth  downward  to  a  flame,  which  should  be  a 
foot  or  two  removed,  and  the  H  will  light  with  a  slight  noise  if  pure. 
Carry  the  tube  back  and  lower  it  for  an  instant  over  the  delivery  tube. 
Hold  the  flame  inside  the  mouth  of  a  dry  bottle  and  continue  until  the 
liquid  deposited  collects  in  drops.  Taste  it  and  test  it  with  litmus  pa- 
per. What  is  it?  In  case  the  evolution  of  gas  becomes  too  slow  it  is 
due  to  exhaustion  of  the  acid  and  the  accumulation  of  zinc  sulfate,  or 
chloride  if  hydrochloric  acid  was  used.  Merely  adding  more  acid 
does  not  suffice.'  It  is  better  to  pour  off  the  solution  and  replace  it 
with  fresh  acid  as  in  the  beginning. 

Fill  a  jar  by  pouring  hydrogen  upward  into  it  from  another  jar 
and  prove  the  presence  of  H  in  the  jar  by  applying  a  flame  to  its 
mouth. 

Hold  a  jar  of  H  mouth  downward,  remove  cover  and 
insert  a  bit  of  lighted  candle  held  on  the  end  of  a  file  or  wire.  Remove 


10 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


candle  when  it  will  light  again.  Repeat  several  times.  Does  a  candle 
burn  in  hydrogen? 

When  the  evolution  of  gas  in  the  flask  has  nearly  ceased  filter 
the  liquid  into  a  dish.  If  sulfuric  acid  was  used  crystals  of  zinc"  sul- 
fate  will  appear  when  it  becomes  cold,  or  after  it  has  been  evaporated 
down  one  half  and  cooled  again.  If  hydrochloric  acid  was  used  eva- 
porate to  dryness  to  obtain  zinc  chloride.  Place  some  of  it  on  a  glass 
plate  and  observe  at  the  next  period.  Is  it  hygroscopic? 

13. — Hydrogen  from  Water:  (a)  Sodium  should  be  handled  with  the 
rorceps  or  dry  hands.  Press  a  piece  of  the  metal  tightly  into  a  22  car- 
cridge  shell.  Fill  a  test  tube  with  water,  drop  the  shell  into  the  trough 
d,nd  collect  the  gas  by  displacement  of  water.  Test  the  gas  for  hydro- 
gen. 

(b)  Set  up  the  apparatus  as  shown  in  fig.  4  using  a  half  inch  gas 
pipe  at  least  a  foot  long.  Put  near  the  middle  of  tube  about  10  grams 
iron  turnings  held  in  place  by  loose  plugs  of  steel  wool.  To  protect 
the  stoppers  wrap  cotton  around  the  ends  of  the  tube  and  keep  it  sat- 
urated with  water.  Use  a  very  small  flask  and  only  lOc.c.  of  water  in 
it.  Heat  strongly  the  middle  of  the  iron  tube  and  after  several  min- 
utes boil  the  water  in  flask.  After  the  air  has  been  driven  out  of  the 
tube  collect  and  test  the  hydrogen.  Prevent  a  back  flow  by  removing 
delivery  tube  from  the  trough  before  removing  burner  from  the  flask. 


Fig.  4. 

The  experiment  may  be  continued  till  a  large  amount  of  the  iron  is 
oxidized  when  it  may  be  used  in  14b;  or,  when  cold  more  of  the  iron 
may  be  "rusted"  by  wetting  it  and  putting  aside  for  a  day  or  two.  It 
may  then  be  heated  and  dried  with  a  current  of  air  before  use  in  14b. 

In  this    experiment  the  reaction  is  shown  from  left  to  right: 

3Fe-f4H2O=(reversibly)Fe3O4+4H2 

while  in  14b  it  goes  from  right  to  left,  if  the  contents  of  this  tube  or 
magnetite  is  there  used  as  provided  for. 

14. — Hydrogen  as  a  Reducing  Agent:    The  experiment  may  be  per- 
formed as  in  (a)  or  (b). 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


II 


(a)  Place  in  a  piece  of  hard  glass 
tubing  a  column  of  granulated  cop- 
per oxide  using  loose  plugs  of  as- 
bestos to  hold  the  oxide  in  place. 
The  tube  may  be  drawn  out  and 
bent  at  right  angle  as  shown  in 
Fig.  5  or  a  small  right  angled  tube 
may  be  connected  with  its  outer 
end  by  rubber.  Connect  with  the 
flask  which  contains  zinc,  and  add 
dilute  sulphuric  acid.  Turn  the 
small  tube  upward  and  determine 
when  the  air  is  wholly  expelled  by 
lighting  the  hydrogen  with  a  test 

tube  as  in  Ex.  12.  Now  turn  the  tube  downward,  place  it  in  a  test  tube 
surrounded  with  water  and  heat  the  copper  oxide,  beginning  at  the  end 
of  the  column  nearer  the  flask.  Continue  till  the  copper  oxide  is  all 
reduced,  that  is,  all  changed  to  red  copper.  Examine  the  liquid  that 
has  collected  in  the  test  tube.  What  is  it? 

(b)  For  the  glass  tube  and  copper  oxide  in  (a)  substitute  the 
iron  tube  and  contents  from  (13b)  or  use  the  iron  tube  and  about  5 
grams  of  ferric  oxide.  First  fill  the  tube  completely  with  hydrogen  as 
proved  by  lighting  it  with  a  test  tube,  heat  the  oxide  strongly,  protect- 
ing the  stoppers  with  the  wet  cotton  as  stated  in  13b. 

With  a  moderate  flow  of  hydrogen  continue  the  reduction  10  to  15 
minutes.  What  is  the  source  of  the  water  in  the  test  tube?  When 
cold  and  if  ferric  oxide  was  used  test  some  of  the  contents  of  the  tube 
in  con.  HC1  warming.  What  gas  is  evolved?  What  evidence  have  you 
that  the  reaction  in  13b  is  reversed?  What  is  a  reversible  reaction? 
What  governs  the  direction  of  this  one?  How  could  you  produce  a 
state  of  equilibrium  in  it?  What  is  the  meaning  of  equilibrium? 

WATER. 

15. — Solubility  of  Solids:  Some  substances  are  very  soluble  in 
cold,  still  more  soluble  in  hot  water.  Shake  3  grams  of  powdered  so- 
dium nitrate  or  ammonium  chloride  with  5c.c.  of  water  in  a  test  tube 
till  all  is  dissolved.  Is  there  any  change  in  temperature  of  the  water? 
Now  add  3  grams  more  and  shake.  Shake  and  determine  whether  it 
all  dissolves.  If  not,  heat  till  dissolved.  What  occurs  on  cooling  to 
room  temperature? 

Some  substances  are  very  soluble  in  hot  water,  slightly  in  cold 
water,  and  such  may  easily  be  purified  by  crystallization  from  hot  so- 
lution. By  heating  dissolve  4  grams  of  ordinary  potassium  chlorate 
in  lOc.c.  distilled  water.  Filter  boiling  hot  if  there  is  any  turbidity  or 


12  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

solid  matter  suspended  in  the  liquid.  When  the  solution  is  cold  filter 
off  the  crystals,  setting  filtrate  aside  and  wash  crystals  with  a  little 
cold  water,  dissolve  a  portion  of  them  in  distilled  water  and  add  sil- 
ver nitrate  to  the  solution  and  to  filtrate.  Compare  the  amount  of 
precipitate  obtained  with  silver  nitrate  in  the  two  tubes.  It  is  due  to 
a  chloride,  commonly  found  in  ordinary  potassium  chlorate. 

Some  substances  are  little  more  soluble  in  hot  water  than  in  cold 
wau-r.  Shake  5  grams  of  common  salt  with  lOc.c.  of  water  in  a  test 
tube.  Note  the  amount  of  salt  undissolved.  Boil  the  solution  for  a  few 
moments  to  saturate  the  water  with  salt.  Note  again, the  amount  of 
salt.  Filter  the  solution  boiling  hot  and  cool  to  room  temperature. 
Does  much  salt  crystallize  out?  Why  not?  Compare  the  solubility  01 
salt  in  hot  and  cold  water  with  that  of  potassium  chlorate.  A  few  sub- 
stances are  even  less  soluble  in  hot  water  than  in  cold.  Examples  are 
calcium  sulfate  (gypsum)  and  slaked  lime. 

16.  Solubility  of  Liquids  in  Liquids,  and  the  Separation  of  Solutes 
between  two  non-miseible  Solvents:  Measure  accurately  in  a  cylinder 
about  50c.c.  of  water  reading  at  the  lower  surface  of  the  meniscus.  Now 
carefully  pipet  into  the  cylinder  25c.c.  of  common  alcohol.  Mix  thor- 
oughly and  read  the  volume.  Is  the  volume  now  the  sum  of  those  of 
the  alcohol  and  water?  Do  water  and  alcohol  dissolve  each  other 
completely? 

In  a  test  tube  place  about  lOc.c.  of  water  and  about  2c.c.  of  chloro- 
form and  shake.  Let  stand  and  note  whether  they  are  mixed.  Add  a 
crystal  of  iodine  and  shake  for  some  time  and  let  stand.  Which  li- 
quid takes  up  most  of  the  iodine? 

As  above  try  to  mix  lOc.c.  of  \vater  and  2c.c.  of  carbon  disulfide. 
Add  about  oc.c.  of  bromine  water  and  shake.  Which  liquid  takes  up 
most  of  the  bromine? 

To  another  tube  add  water  and  a  little  benzene,  and  a  small  crystal 
of  potassium  permanganate  and  shake.  Which  liquid  takes  up  the 
permanganate? 

Determine  whether  alcohol  and  chloroform,  benzene  and  chloro- 
form, carbon xdisulfide  and  benzene  will  mutually  dissolve  each  other. 

17— Solubility  of  Gases:  Gases  also  vary  widely  in  their  solu- 
bility in  water.  In  text-book  see  solubility  of  oxygen,  nitrogen  and  hy- 
drogen sulfide  which  do  not  act  chemically  with  the  water,  dissolve 
only  moderately  and  obey  Henry's  law,  (which  see).  Others  which 
generally  act  chemically  with  water  dissolve  in  very  large  amounts. 

In  a  test  tube  or  flask  heat  hydrant  water  and  observe  gas  bub- 
bles given  off  before  the  water  begins  to  boil.  Try  the  same  with  dis- 
tilled water.  Does  it  contain  dissolved  gases.  Why  does  water  when 
drawn  from  the  hydrant  sometimes  look  milky  and  quickly  become 
clear  on  standing? 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  13 

Note  carefully  the  odor  from  a  solution  of  ammonia  due  to  am- 
monia given  off.  Pour  about  Ic.c.  of  the  solution  into  lOc.c.  of  water 
in  a  dish  and  test  with  turmeric  paper.  Boil  till  the  water  is  half 
evaporated  and  test  again  with  the  paper.  What  is  the  effect  of  heat 
on  the  solubility  of  gases.  To  a  degree  the  following  shows  one  of 
many  exceptions  to  the  general  rule: 

Boil  lOc.c.  of  dilute  hydrochloric  acid  till  three-fourths  cf  it  has 
evaporated.  Boil  the  same  volume  of  concentrated  hydrochloric  acid 
till  one  half  has  evaporated.  Add  to  each  sample  of  boiled  acid  wheji 
cold,  the  same  amount  of  granulated  zinc.  Do  they  seem  to  act  on  zinc 
at  about  the  same  rate?  A  more  accurate  determination  would  show 
that  they  have  the  same  concentration.  By  boiling  long  enough 
samples  of  dilute  and  concentrated  hydrochloric  acids  one  arrives  at 
the  same  result;  namely  acids  of  concentration  20.2  per  cent. 

18. — Chemical  Action  of  Water:  Refer  back  to  13  ,  for  the  action 
of  water  on  sodium  and  on  iron,  and  to  9  for  its  action  on  sodium  di- 
oxide. 

Upon  a  piece  of  quick  lime  drop  water  slowly  until  the  water  is  no 
longer  absorbed  and  the  piece  of  lime  looks  wet.  Place  it  in  a  dish 
and  note  what  occurs.  This  is  the  familiar  "slaking"  of  lime. 

10.— Water  in  Combination:  In  a  test  tube  heat  a  small  amount  of 
copper  sulfate,  observing  water  given  off  and  changed  appearance.  In 
the  same  way  try  borax,  alum,  sodium  phosphate.  These  are  "hyd- 
rates" and  the  water  they  contain  is  called  "water  of  crystallization." 
For  contrast  try  potassium  sulfate  and  common  salt. 

In  a  weighed  porcelain  crucible  with  lid  weigh  accurately  about  4 
g.  of  barium  chloride,  place  the  crucible  on  a  triangle  over  a  burner 
and  heat  ten  minutes.  When  cold  weigh  the  crucible  and  contents 
again  and  from  the  first  weight  of  the  barium  chloride  and  the  loss  on 
heating,  calculate  the  per  cent,  of  water. 

Efflorescence:  On  a  glass  plate  or  watch-glass  place  a  few  crys- 
tals of  sodic  sulphate,  expose  to  air  till  the  next  laboratory  period. 

Deliquescence:  Expose  to  air  in  dishes  small  pieces  of  calcium 
chloride  and  caustic  potash  till  the  next  laboratory  period  and  record 
results. 

20.  Electrolytic  Decomposition  of  Water:  Support  the  U  tube 
shown  in  fig.  14  with  a  clamp  and  use  the  current  described  in  53. 
Connect  the  side-arms  of  the  tube  with  the  trough  of  water  by  means 
of  short  delivery  tubes.  Fill  the  U  tube  nearly  to  the  side  arms  with 
a  5th  normal  solution  of  sodium  sulfate  already  made  up  and  add  a  few 
drops  of  litmus  solution  to  each  side  of  U  tube.  Start  the  current  and 
at  the  same  time  place  over  the  ends  of  the  delivery  tubes  two  test 
tubes  the  same  size  and  full  of  water,  thus  collecting  the  hydrogen  and 
oxygen  set  free.  Continue  till  the  smaller  volume  equals  about  5c.c. 


14 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


and  compare  volumes.    Note  color  of  the  solution  at  each  electrode  and 
state  cause.    Name  electrodes.    From  which  comes  the  0,  and  the  H? 

21.  Purification  of  Water  by  Distillation:  To  a  small  volume  of 
hydrant  water  add  a  few  drops  of  barium  chloride  solution.  The  white 
precipitate  shows  carbonate  and  sulfate  radicals  present.  Now  add 
dilute  hydrochloric  acid.  The  precipitate  remaining  shows  sulfate 
radical.  To  another  portion  add  a  few  drops  of  dilute  nitric  acid  and 
a  few  of  silver  nitrate.  A  white  precipitate  shows  chloride  present. 

Set  up  apparatus  shown  in  fig.  6,  fill  flask  one-third  full  of  water, 
distill  enough  to  clean  the  tubes  and  reject  it.  Distill  about  lOc.c.  of 
the  water  and  test  it  for  sulfate  and  chloride  radicals.  None  should 
be  obtained,  all  ordinary  mineral  matter  being  left  in  the  boiler  in  dis- 
tillation. 

Try  the  action  of  ammonia  and  so- 
dium hydroxide  on  turmeric  paper 
and  dilute  siilfurio  aoi^  on  blue  lit- 
mus papei  To  one-fourth-  of 
flask  full  of  water  add  about  5c.c.  of 
ammonia,  distill  and  determine  wheth- 
er any  ammonia  distills  over.  Clean 
apparatus  carefully  and  try  sodium 
hydroxide,  adding  a  few  c.c.  to  one- 
fourth  flask  of  water.  Distill  from 
a  fresh  portion  of  water  containing 
a  little  sulfuric  acid  and  a  little 
potassium  permanganate.  Do  they  go 
Fig.  6.  over? 

HYDROGEN  DIOXIDE  (PEROXIDE). 

22. — Preparation:  Place  a  beaker  with  lOOc.c.  of  water  in  water, 
preferably  ice  cold.  Stir  in  little  by  little  5  grams  of  sodium  dioxide, 
Na»O2  the  operation  taking  about  5  minutes.  Even  then  some  of  the 
peroxide  will  be  decomposed  giving  off  oxygen.  Now  add  gradually 
in  the  same  way  dilute  hydrochloric  acid  till  a  drop  of  the  liquid  taken 
out  with  the  stirring  rod  just  turns  blue  litmus  paper  red.  This  gives 
a  solution  of  hydrogen  dioxide,  but  it  contains  also  common  salt.  It 
may  be  used  where  hydrogen  peroxide  is  required  below  save  where  it 
is  used  with  silver  and  lead.  There  the  commercial  peroxide  should  be 
used. 

Rub  in  a  mortar  with  a  little  water  about  half  a  gram  of  starch, 
transfer  to  a  dish  or  beaker,  add  about  lOOc.c.  of  water  and  a  few  crys- 
tals of  potassium  iodide  and  heat  to  boiling.  This  is  known  as  "starch- 
iodide  solution"  and  filter  paper  wet  with  it  is  called  "starch-iodide  pa- 
per." Each  is  frequently  required.  To  a  part  of  the  solution  add  a 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  15 

few  drops  of  the  solution  of  hydrogen  peroxide  which  sets  free  iodine 
and  this  colors  the  starch  blue.  To  see  the  color  by  transmitted  light 
dilute  with  much  water.  This  is  used  as  a  test  for  iodine  or  hydrogen 
peroxide  or  starch.  That  is,  two  being  known  to  be  present  the  pres- 
ence or  absence  of  the  third  can  be  determined  by  the  test. 

23. — Oxidation  with  Hydrogen  Peroxide:  Moisten  a  strip  of  filter 
paper  with  very  dilute  lead  acetate  and  expose  it  to  hydrogen  sulfide 
a  little  of  which  may  be  made  in  a  test  tube  as  in  68.  The  black  sub- 
stance is  lead  sulfide,  PbS.  Pour  upon  the  paper  a  few  drops -of  hy- 
drogen dioxide,  which  will  change  the  lead  sulfide  to  white  lead  sul- 
fate,  PbSO*.  Why  is  this  called  oxidation? 

To  a  solution  of  silver  nitrate  add  NaOH  and  then  carefully  add 
just  enough  of  a  solution  of  ammonium  hydroxide  to  dissolve  the  pre- 
cipitate at  first  formed.  Now  add  commercial  hydrogen  peroxide.  The 
gray  precipitate  is  finely  divided  metallic  silver,  and  the  escaping  gas 
is  oxygen  which  may  be  tested  by  trying  in  the  tube  a  glowing  splint- 
er. This  action  appears  to  be  one  of  reduction,  but  it  is  not  primarily. 
Probably  a  higher  oxide  of  silver  is  formed  and  at  once  decomposes 
into  silver  and  oxygen.  The  next  two  cases  are  of  the  same  sort. 

To  about  a  gram  of  manganese  dioxide  in  a  test  tube  add  hydrogen 
dioxide  and  test  the  gas  with  a  glowing  splinter.  Repeat  using  a  con- 
centrated solution  or  a  few  crystals  of  potassium  permanganate  in- 
stead of  the  manganese  dioxide. 

CHLORIDE. 

24.— Preparation  of  Chlorine:  Chlorine  is  dangerous  if  breathed. 
Experiments  with  it  should  be  conducted  in  hoods.  If  carried  out  on  the 
students'  desks  only  half  the  usual  number  should  work  in  the  room  at 
one  time  and  windows  should  be  freely  opened.  They  should  stand  to 
windward  of  the  apparatus  when  collecting  the  gas.  When  through 
collecting  at  once  place  the  delivery  tube  in  a  test  tube  nearly  full  of 
concentrated  sodium  hydroxide,  and  remove  the  source  of  heat.  Be- 
fore beginning  the  experiments  students  should  read  them  through, 
the  right  plan  in  any  case,  provide  everything  necessary  and  thus  re- 
duce the  time  to  the  minimum.  Make  a  solution  of  starch-iodide  as  di- 
directed  in  22.  , 

Chlorine  is  made  commercially  by  the  electrolysis  of  fused  com- 
mon salt  or  a  solution  of  common  salt.  All  other  methods  depend  up- 
on the  oxidation  of  the  hydrogen  of  hydrochloric  acid.  Several 
methods  may  be  illustrated  on  a  very  small  scale. 


16 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


Upon  a  few  crystals  of  potas- 
sium permanganate  in  a  test  tube 
pour  about  Ic.c.  of  con.  hydrochlo- 
ric acid  and  pour  a  little  of  the 
heavy,  greenish  yellow  gas  into  a 
little  of  the  starch  solution  in  an- 
other tube.  In  the  same  way  treat 
a  little  potassium  dichromate  heat 
and  test  for  chlorine.  Try  in  the 
same  way  lead  dioxide  and  con. 
hydrochloric  acid.  Use  a  very 
little  potassium  chlorate  and  con- 
centrated hydrochlpric  acid,  HC1; 
also,  about  1  cc.  con.  HC1  and  a 
few  drops  of  con.  nitric  acid. 


Fig.  7. 

25.— Set  up  the  apparatus  as  in  fig.  7,  with  the  water  bath  one- 
fourth  full  of  water.  A  copper  can  is  best,  but  a  tin  can  or  a  beaker  will 
serve  for  a  water  bath.  See  that  the  thistle  tube  reaches  nearly  to  the 
bottom  of  the  flask.  In  the  flask  place  25  grams  of  manganese  diox- 
ide MnO2,  and  add  40c.c.  of  concentrated  HC1  diluted  with  lOc.C.  of 
water.  Heat  the  flask  and  collect  five  jars  or  bottles  of  chlorine.  Tlie 
green  color  will  show  when  they  are  full.  They  should  be  well  filled, 
but  a  large  excess  should  not  over-flow  into  the  room.  One  jar  should 
have  the  bottom  wet  with  con.  sulfuric  acid  to  dry  the  gas  for  use  in 
bleaching. 

When  through  collecting  place  the  delivery  tube  in  NaOH  in  a 
test  tube.  Let  the  flask  cool,  removing  water  bath,  and  proceed  with  the 
next  experiment. 

26. — Properties  of  Chlorine:  To  show  bleaching  action  and  the 
need  of  water,  suspend  in  the  jar  of  dry  chlorine  strips  of  colored 
cotton  cloth  and  litmus  paper.  After  a  few  moments  note  any  fading 
of  the  colors,  then  moisten  the  strips  and  suspend  again  in  the  jar, 
and  note  effect.  Refer  to  a  text-book  for  information  on  hypochlorous 
acid  and  bleaching  with  chlorine.  What  does  the  bleaching? 

Into  a  jar  of  chlorine  pour  successively  with  shaking,  small  vol- 
umes of  solutions  of  litmus,  cochineal,  much  diluted  ink. 

From  a  piece  of  antimony  scrape  with  a  knife  a  very  little  of  the 
metal  letting  it  fall  into  a  jar  of  chlorine.  Using  forceps  or  tongs  heat 
to  redness  a  strip  of  copper  foil  and  lower  it  into  the  same  jar  of  chlo- 
rine. If  "Dutch  metal"  is  used  it  need  not  be  heated. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  17 

Burn  the  laboratory  gas  at  the  end  of  a  glass  tube  and  lower  the 
small  flame  into  a  jar  of  chlorine.  What  is  the  black  substance?  When 
the  green  color  has  disappeared  blow  breath  over  mouth  of  jar,  which 
will  give  a  fog,  consisting  of  droplets  of  water  containing  HC1.  Will 
carbon  burn  in  Cl?  Try  an  ignited  piece  of  charcoal. 

On  a  deflagration  spoon  lower  a  small  bit  of  white  phosphorus  in- 
to a  jar  of  Cl  and  avoid  inhaling  the  Cl  or  fumes  of  PC13.  It  should 
soon  melt,  then  take  fire. 

Now  take  the  flask  and  all  jars  to  the  sink  best  under  hood  and 
standing  well  back  fill  them  with  water.  Wash  well  any  remaining 
MnOi.  and  place  it  in  a  vessel  provided  for  that  purpose. 

HYDROCHLORIC  ACID. 

27. — Preparation :  Read  the  experiment  through  and  have  all  the 
necessary  materials  at  hand  so  that  once  begun  the  experiment  may 
be  carried  through  rapidly.  ^ 

All  soluble  chlorides  give  hydrochloric  acid  when  treated  with 
concentrated  sulfuric  acid.  For  many  reasons  sodium  chloride,  com- 
mon salt,  is  to  be  preferred. 

Set  up  the  apparatus  as  in  fig.  7,  omitting  the  water  bath.  In  the 
flask  place  20  grams  of  sodium  chloride.  Dilute  35  grams  (20c.c.)  of 
concentrated  sulfuric  acid  by  pouring  it  slowly  with  stirring  into  8c.c, 
of  water  in  a  beaker.  Pour  slowly  into  the  flask  and  let  stand  a  few 
moments  till  acid  and  salt  are  in  contact  throughout  then  apply  a  low 
heat,  best  using  a  burner  with  crown  top.  Collect  the  gas  in  dry  jars 
or  bottles  in  the  same  way  as  chlorine.  Abundant  fumes  will  indicate 
when  the  jars  are  full.  After  collecting  two  jars,  fill  a  dry  bottle 
which  has  been  fitted  with  a  stopper  and  a  short  piece  of  tubing  one 
end  flush  with  the  large  end  of  stopper  and  the  other  drawn  out  and 
cut  off  so  as  to  leave  a  small  orifice.  It  should  reach  at  least  to  the 
middle  of  the  bottle's  length.  Insert  stopper  when  full  and  place  the 
bottle  mouth  downward  in  water.  Press  down  the  bottle  and  pour  cold 
water  over  it  to  start  the  absorption.  A  fountain  will  result. 

Into  a  test-tube  of  water  insert  the  delivery-  tube  so  that  it  reaches 
a  very  little  below  the  surface  of  the  water.  Note  ready  absorption  of 
the  gas.  Is  the  solution  lighter  or  heavier  than  water?  Lower  the 
tube  as  necessary  to  absorb  all  the  gas.  Why  does  the  water  become 
warm?  Continue  till  the  water  is  nearly  saturated,  then  try  the  action 
of  small  portions  of  the  solution  on  a  little  zinc,  marble,  sodium  car- 
bonate. Compare  its  action  with  that  of  the  dilute  HC1  from  the  shelf. 
Test  its  action  on  blue  litmus  paper.  To  a  little  of  the  solution  add  a 
few  drops  of  silver  nitrate  then  a  few  drops  of  dilute  nitric  acid.  The 
precipitate  is  silver  chloride.  This  is  a  test  for  hydrochloric  acid  or 
a  chloride.  Repeat  using  instead  of  HC1  a  few  drops  of  a  solution  of 


18  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

salt  or  of  any  other  chloride.    With  a  known  chloride  could  it  be  used 
as  a  test  for  silver?    Save  flask  and  its  contents  for  29. 

28.— Acid,  Base,  Salt,  Neutralization:  Wet  the  inside  of  a  jar  with 
a  solution  of  concentrated  ammonia  and  pour  out  excess  of  liquid. 
Cover  jar  with  a  glass  plate.  Place^  it  over  a  jar  of  the  HC1  gas  from 
the  last  experiment,  bringing  the  jars  mouth  to  mouth  and  remove 
plate.  The  ammonia  unites  with  acid  forming  ammonium  chloride,  a 
salt. 

Standing  well  away  drop  one  or  two  small  bits  of  sodium  into  a 
little  distilled  water  in  a  bottle.  When  the  action  is  over  test  the  wa- 
ter with  turmeric  paper.  The  action  of  sodium  on  water  gives  sodium 
hydroxide,  a  base,  and  what  gas?  (see  13). 

Test  sodkim  hydroxide  from  shelf  bottle  with  the  papers.  Pour 
about  5c.c.  of  the  alkali  into  a  dish  and  add  dilute  hydrochloric  acid  till 
the  solution  turns  litmus  paj>er  red,  testing  by  taking  out  a  drop  of  the 
solution  with  a  stirring  rod  and  touching  the  paper ;  never  place  pa- 
pers in  the  solution  or  dip  them  into  it.  Now  add  a  few  drops  of  NaOH 
or  dilute  HC1  as  may  be  necessary  with  stirring,  till  the  solution 
changes  neither  turmeric  nor  blue  litmus  paper.  It  is  now  neutral. 
Evaporate  to  dryness,  taste  the  residue.  Place  upon  it  a  few  drops  of 
con.  sulfuric  acid  and  note  odor  of  the  gas.  What  was  the  solid  resi- 
due? 

29. — The  preparation  of  HC1  from  salt  and  sulfuric  acid  gives  a 
good  example  of  a  reversible  reaction: 

NaCl+H2SO4=  (reversibly)  HNaSO4+HCl. 

HC1  is  a  stronger  acid  than  H2SO4.  In  the  cold  or  in  a  small  closed 
space  even  on  heating  the  reaction  would  not  complete  itself  to  the 
right.  But,  the  HC1  is  easily  volatile  and  heating  drives  it  out  of 
solution  and  away  so  that  it  cannot  react  with  HNaSO4  to  the  left.  Sul- 
furic acid  is  volatile  only  at  very  high  temperature.  Try  restoring 
HC1  thus:  Dissolve  with  the  least  volume  of  water  and  heating,  the 
contents  of  the  flask  used  in  preparing  HC1.  Cool  some  of  the  solution 
of  HNaSO4  in  a  test  tube  and  add  a  few  c.c.  of  con.  HC1  from  shelf, 
which  will  precipitate  sodium  chloride.  Why  is  sulfuric  acid  used  to 
prepare  easily  volatile  acids  from  their  salts?  Why  are  the  reactions 
completed  when  heat  is  used? 

30.— Preparation  of  Bleaching  Powder,  Potassium  Hypochlorite 
and  Potassium  Chlorate:  Study  these  subjects  in  a  text-book  and  read 
the  experiment  through. 

Arrange  the  apparatus  as  shown  in  fig.  8.  The  wash  bottle  con- 
tains diluted  sulfuric  acid,  made  by  adding  40c.c.  of  the  con.  acid  with 
stirring,  to  lOc.c.  of  water.  Its  purpose  is  to  remove  the  greater  part 
of  the  HC1  and  water  from  the  Cl.  The  horizontal  test  tube  contains 
about  2  grams  of  dry  slaked  lime  spread  evenly  throughout  irs 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


19 


O 


length.  The  upright  test  tube  should  contain  a  solution  of  4  grams  of 
potassium  hydroxide  dissolved  in  12c.c.  of  water.  It  must  be  cooled 
and  kept  cold  by  surrounding  with  water  as  shown. 

Charge  the  flask  with  MnO2  and  diluted  HC1  as  in  25.  Maintain  a 
moderate  stream  of  chlorine  for  about  15  minutes  or  until  the  larger 

portion  of  it  seems  to  pass 
through  the  solution  in  the  second 
test  tube.  Now  preserve  half  the 
contents  of  this  tube  as  potassium 
hypochlorite,  and  heat  the  re- 

j£,  |>|  mainder    to    boiling    and    without 

(  ™  cooling  continue  to  pass  chlorine 

into  it  for  about  five  minutes  or 
till  a  drop  taken  out  with  a  stir- 
ring rod  does  not  feel  soapy  to  the 
fingers.  Now  heat  to  boiling  and 
filter  the  solution.  Cool  by  stand- 
ing the  tube  in  water,  when  crys- 
tals of  KC1O3  form.  When  quite 
cold  filter  them  off  and  wash  with 
a  little  cold  water.  Test  the  ni- 
trate with  silver  nitrate.  What 
was  formed  *  besides  potassium 
chlorate?  Dissolve  the  chlorate 
Fig.  8.  by  passing  about  3  cc.  of  boil- 

ing water  through  the  filter  several  times,  cool,  let  crystallize,  pour 
off  the  water  from  the  crystals,  dissolve  them  in  water  and  add 
silver  nitrate.  Compare  the  first  and  second  precipitate  formed  by  sil- 
ver nitrate.  Pure  chlorate  would  give  no  precipitate. 

Into  a  little  of  the  hypochlorite  solution  place  a  drop  of  diluted 
ink,  into  another  .portion  a  bit  of  colored  cloth.  After  noting  any 
bleaching  add  a  little  dilute  acid  to  each  solution  and  note  effect.  To 
a  third  portion  add  a  few  drops  of  strong  solution  of  ammonium  hyd- 
roxide. What  gas  is  given  off? 

Try  the  action  of  dilute  acid  on  a  little  of  the  dry  bleaching  pow- 
der from  the  horizontal  tube.  Dissolve  a  portion  of  it  so  far  as  pos- 
sible, filter  and  try  the  action  of  the  filtrate  on  ink,  colored  cloth,  lit- 
mus paper,  before  and  after  adding  dilute  acid. 

To  show  the  instability  and  oxidizing  power  of  potassium  chlorate 
mix  on  paper  5  grams  of  the  salt  and  5  grams  of  powdered  sugar, 
but  do  not  grind  them  together  in  a  mortar.  Place  the  mixture  on  an 
iron  plate,  take  out  a  drop  of  con.  H2SO4  with  the  stirring  rod,  stand 
well  away  and  drop  the  acid  upon  the  mixture. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


BKOMENE 


IODOE. 


31.  —  Preparation  of  Bromine:  Place  5  grams  of  manganese  diox- 
ide and  5  grams  of  sodium  bromide  on  paper,  hold  the  neck  of  the  re- 
tort pointing  slightly  upward  and  slide  the  mixture  in  at  the  tubulus 
without  letting  it  fall  into  the  neck  of  the  retort.  Support  the  retort 
as  in  fig.  9,  add  50c.c.  of  dilute  sulfuric  acid  through  a  funnel.  The 
flask  should  contain  about  lOOc.c.  of  water  and  the  tip  of  the  retort 
neck  should  dip  under  its  surface.  Apply  heat  and  continue  till  all 
the  bromine  has  distilled  over.  Move  the  flask  away  so  as  to  bring 
the  neck  of  the  retort  above  the  surface  and  then  remove  the  burner. 
Is  bromine  soluble  in  water?  Is  it  heavier  than  water?  Set  aside  the 

flask  containing  bromine  for  lat- 
er use. 

Give  two  reasons  why  we  do 
not  prepare  Br  in  the  same  way 
as  Cl;  that  is,  by  the  action  of 
hydrobromic  acid  on  manganese 
dioxide? 

May  Cl  be  prepared  by  the 
same  method  used  for  Br;  that 
is,  by  use  of  NaCl.  MnO2  and  di- 
luted H2SO4?  Try  it  in  a  small 
way  in  a  test  tube,  but  strength- 
en the  acid  by  adding  about  Ic.c. 
con.  H,S04  to  2c.c.  of  dilute  acid. 
Write  equations  for  preparation 
of  eaph. 

~~     Why  is  it  not  best  to  use  con. 

Fig.  9.  sulfuric  acid  in  preparing  either 

Br  or  Cl  by  this  method? 

31.  By  the  same  method  prepare  a  little  iodine,  using  about  O.r> 
gram  of  sodium  iodide,  a  gram  of  MnO2  and  5c.c.  of  a  mixture  of  di- 
lute and  con.  suifuric  acid.  Note  sublimed  iodine  near  the  mouth 
of  the  test  tube.  Compare  the  reaction  with  that  in  the  preparation 
of  Cl  and  Br  by  the  same  method. 

Heat  a  few  crystals  of  iodine  in  a  dry  test  tube.  Does  it  form 
a  liquid  before  subliming?  What  is  sublimation?  Note  crystals 
higher  up  on  walls  of  the  tube. 

33.—  Hydriodic  Acid  and  Comparison  of  HCI,  HHr,  HI:  Treat 
very  small  amounts  of  sodium  chloride,  sodium  bromide  and  sodium 
iodide  with  a  few  dropi  of  con.  sulfuric  acid.  How  did  you  prepare 
HCI?  Could  HBr  be  prepared  in  the  same  way?  What  is  the  black 
substance  set  free  by  the  action  of  the  acid  on  sodium  iodide? 


LABORATORY  MANUx\L  OF  GENERAL  CHEMISTRY  21 

When  salt  and  sulfuric  acid  are  heated  together  hydrochloric 
acid  is  set  free,  but  no  chlorine.  Though  H2SO4  may  be  reduced  and 
HC1  oxidized,  they  are  too  stable  to  act  on  each  other. 

Though  statements  to  the  contrary  are  common,  hydrobromic 
acid  may  be  prepared  in  precisely  the  same  way  as  you  prepared 
HC1,  remembering  that  the  sulfuric  acid  there  used  was  somewhat 
diluted.  A  very  little  HBr  is  oxidized  by  the  H2SO4  giving  a  trace 
of  free  bromine  and  sulfur  dioxide.  Write  the  equation.  But  when 
the  gas  is  absorbed  with  water  and  the  solution  is  distilled,  the 
trace  of  Br  quickly  passes  off,  and  the  sulfur  and  its  compounds 
that  may  be  liberated  are  oxidized  to  sulfuric  acid  which  remains 
to  the  last  in  the  distilling  vess.el. 

Hydriodic  acid  cannot  be  prepared  by  this  method,  since  the 
HI  is  almost  completely  oxidized  to  free  iodine  and  water,  and  the 
sulfuric  acid  is  reduced  to  sulfurous  acid  and  even  to  hydrogen  sul- 
fide. 

Compare  the  degrees  of  stability  of  HC1,  HBr,  HI,  and  compare 
these  with  their  heats  of  formation  by  consulting  a  reference  book. 

34. — Preparation  of  Hydriodic  Acid:  Arrange  a  test  tube  as  in 
fig.  10  high  enough  to  permit  heating  and  having  a  right  angled  de- 
livery tube.  Place  in  the  dry  test  tube  5  grams  powdered  iodine  and 
on  top  of  it  0.5  red  phosphorous,  shake  to  mix,  put  tube  in  place  and 
warm  till  the  P  and  I  react.  When  cool  place  delivery  tube  in 
small  flask  or  bottle  and  by  means  of  the  pinch  cock  drop  into  test 
tube  about  20  drops  of  water.  Warm  gently  and  collect  the  receiver 
full  of  HI  and  stopper  it.  Now  place  the  delivery  tube  in  water  in 
test  tube  to  absorb  the  gas,  but  have  the  tip  of  delivery  tube  just 
above  the  surface  of  the  water. 

Prepare  a  little  Cl  in  a  test  tube  by  warming  a  few  crystals  of 
KC1O:;  and  a  little  con.  HC1.  Pour  some  of  the  Cl  into  the  flask  con- 
taining the  HI.  What  is  set  free? 

Test  the  solution  of  HI  in  water  with  silver  nitrate.  What  was 
the  action  of  the  P  and  I?  What  was  the  action  of  this  compound 
and  the  water?  What  is  hydrolysis.  Write  all  equations. 

This  same  method  with  some  modifications  is  often  used  for 
the  preparation  of  hydrobromic  acid.  (See  the  text  book  for  de- 
scription.) 

35. — Powder  a  few  crystals  of  iodine  in  a  mortar,  place  it 
in  a  test  tube  half  full  of  water,  and  pass  into  the  tube  hydrogen 
sulfide  gas  prepared  as  in  68.  Have  the  delivery  tube  reach  quite 
to  the  bottom  of  the  test  tube.  When  all  the  iodine  has  disappeared 
boil  the  solution  for  a  time  to  expel  the  excess  of  hydrogen  sulfide 
and  filter,  several  times  if  necessary  to  get  rid  of  all  the  sulfur. 


22  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

Test  the  filtrate  with  blue  litmus  paper.  Set  the  solution  aside  f.o 
be  used  in  the  next  experiment.  Hydrogen  sulfide  here  acts  as  a 
reducing  agent.  Write  the  equation. 

:W.— Properties  of  Chlorine,  Bromine,  Iodine  and  Their  Com- 
pounds: Try  to  dissolve  a  little  powdered  iodine  in  water.  Now 
pour  off  most  of  the  water,  add  a  few  crystals  of  sodium  iodide  and 
shake.  It  is  supposed  the  solution  now-  contains  NaI3. 

Try  the  solubility  of  I  by  pouring  about  Ic.c.  of  carbon  disul- 
fide  upon  a  crystal  of  I  in  a  test  tube  and  shaking.  Now  add  water 
to  the  tube,  shake  and  let  carbon  disulfide  settle.  In  which  liquid 
is  most  of  the  iodine?  Try  the  solubility  of  iodine  also  in  alcohol 
and  in  chloroform.  Save  the  solutions  of  iodine. 

Make  a  solution  of  •  starch  by  boiling  about  half  a  gram  of 
starch  in  a  dish  or  beaker  half  full  of  water  and  with  stirring. 
Pour  a  little  of  the  starch  solution  into  a  beaker  of  water,  add  a 
little  iodine  solution.  This  is  a  good  test  for  free  iodine.  Add  a 
little  sodium  iodide  solution  and  a  little  of  the  starch  solution  to  a 
bleaker  of  water.  To  one  portion  add  a  little  bromine  water  and  to 
the  other  chlorine  water.  Is  iodine  set  free  by  the  Cl  and  the  Br? 
To  a  solution  of  sodium  bromide  add  chlorine  water.  Is  bromine 
set  free?  To  concentrate  the  bromine  and  also  to  show  its  solu- 
bility, to  the  liquid  in  the  test  tube  add  a  little  carbon  disulfide, 
shake  and  let  the  latter  settle. 

Arrange  the  three  halogens  in  the  order  of  the  ability  of  each 
to  replace  the  others,  and  compare  this  order  with  that  of  the  sta- 
bility of  their  compounds  with  hydrogen. 

To  a  little  solution  of  iodine  and  to  blue  starch-iodine  solution 
add  a  solution  of  sodium  thiosulfate  till  the  colors  disappear. 

To  show  the  reducing  power  of  hydriodic  acid  make  very  di- 
lute solutions  of  potassium  dichromate  and  potassium  permangan- 
ate, add  a  little  dilute  sulfuric  acid  and  then  some  of  your  solution 
of  hydriodic  acid.  The  chromate  and  permanganate  are  reduced 
and  the  solution  is  colored  brown  by  iodine.  For  the  same  purpose 
treat  a  little  of  a  solution  of  sodium  iodate,  NaIO3,  with  dilute  H2S04  and 
add  some  of  your  solution  of  HI.  Refer  to  a  text  book  and  write 
the  equations. 

37. — Tests  for  the  Halogens:  How  may  each  of  the  halogens  in 
the  free  condition  be  recognized  and  tested  for?  (See  36).  The 
following  applies  to  them  in  the  form  of  soluble  halides : 

To  three  tubes  containing  respectively  a  little  dilute  HC1  or  so- 
lution of  any  chloride,  solution  of  any  bromide,  solution  of  any 
iodide,  add  a  few  drops  of  silver  nitrate.  Compare  the  colors  of 
the  three  silver  halides.  Try  to  dissolve  a  little  of  each  with  am- 
thiosulfate.  Try  other  portions  with  dilute  nitric  acid.  Dissolve  a 


LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY  23 

aionium  hydroxide.  Try  other  portions  with  a  solution  of  sodium 
little  of  each  with,  potassium  cyanide,  KCN,  being  careful  not  to  get  it 
on  the  hands. 

38. — Hydrofluoric  Acid:  Slowly  and  evenly  heat  a  glass  plate 
over  a  burner  flame.  Rub  over  the  surface  a  piece  of  bees-wax  so 
as  to  make  a  continuous  film  of  the  melted  wax.  When  the  wax  has 
hardened,  write  on  the  plate  with  a  pointed  file  or  knife,  cutting 
4uite  through  the  wax.  In  a  lead  dish  place  a  little  ammonium 
fluoride,  or  powdered  calcium  fluoride,  moisten  with  con.  sulfuric 
acid.  Cover  the  dish  with  plate,  wax  side  down,  and  apply  a  low 
heat  to  the  dish  so  as  not  to  melt  the  lead.  After  a  few  moments, 
remove  the  wax  and  examine  the  glass. 

To  test  for  a  fluoride  place  a  little  of  it  mixed  with  a  little 
sand,  in  a  dry  test  tube,  drop  upon  the  mixture  a  little  con.  H2SO4 
and  heat.  While  heating  hold  in  the  tube  a  wet  stirring  rod.  Where 
wet  the  rod  will  be  covered  with  a  jelly-like  layer  of  silicic  acid 
Why  cannot  a  solution  of  H2F2  be  kept  in  a  glass  bottle? 

EQUIVALENT  WEIGHTS. 

39. — Equivalent  of  Zinc:  Set  up  the  apparatus  as  shown  in  fig. 
tO  and  prove  that  all  joints  are  tight  by  placing  water  in  the  funnel 
covering  the  end  of  the  delivery  tube  with  the  wet  finger  and  open- 
ing the  pinch  cock.  The  water  should  not  run  down.  The  end  of 
the  delivery  tube  should  be  bent  upward  and  securely  placed  under 
the  jar.  Use  pure,  bright  granulated  zinc. 


Fig.  10. 

Weigh  with  great  care  not  less  than  0.2  grams  nor  more  than 
0.25  grams  of  the  zinc  for  every  lOOc.c.  that  the  jar  or  bottle  will 
hold.  If  an  analytical  balance  is  used  one  weighing  will  suffice. 


^  I  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

Horn  pan  balances  are  not  very  accurate  and  rarely  in  exact  equi- 
librium. In  this  case  weigh  the  zinc  on  one  pan,  then  on  the  other 
and  take  half  the  sum  of  the  weights.  What  errors  in  the  balance 
will  this  practically  eliminate?  Place  zinc  in  the  test  tube,  fill 
collecting  jar  completely  with  water,  and  add  15c.c.  of  con.  HC1  to 
the  funnel.  Let  in  all  the  acid  and  close  the  clamp.  When  all  the 
Zn  is  dissolved  bring  the  level  of  the  water  in  the  jar  to  that  in  the 
trough  and  securely  place  on  the  cover.  Dry  the  outside  of  the  jar 
and  weigh  on  the  platform  balance.  Pill  it  completely  with  water 
and  weigh  again.  The  difference  in  grams  less  the  volume  of  acid 
gives  the  volume  of  the  hydrogen  in  cc.  (V).  Find  the  temperature  of 
the  water  in  the  trough  (t),  and  the  reading  of  the  barometer  (P).  Cal- 
culate the  volume  (V)  to  normal  conditions  by  use  of  the  formula, 
V'=VX273X(P— aq.  tens,  at  t°)-^(273-H°)  X760.  Why?  The  weight 
of  H=V'X. 00009.  The  equivalent  of  zinc  equals  its  weight  divided  by 
the  weight  of  HX  1.008,  and  its  atomic  weight  equals  twice  this  value. 

40. — Equivalents  of  Other  Metals:  Find  the  equivalent  of  one  or 
more  of  the  following,  using  not  more  than  the  weights  given  for  each 
100  c.c.  that  the  receiving  vessel  holds.  Aluminium  0.07  gram;  mag- 
nesium, 0.08;  .iron,  0.15.  In  the  cases  of  aluminium  and  magnesium 
it  is  better  to  place  the  test  tube  in  a  beaker  of  water  to  keep  the  tem- 
perature down,  and  to  use  dilute  acid.  In  the  case  of  iron  it  may  be 
necessary  to  warm  the  con.  HC1  used.  Collect  the  hydrogen  and  pro- 
ceed precisely  as  in  the  previous  experiment. 

41, — Equivalent  of  Chlorine  (a) :  In  normal  times  porcelain  Gooch 
crucibles  are  cheap,  and  if  they  are  available  this  method  should  be 
used;  otherwise  use  (b). 

Weigh  accurately  about  0.5  gram  pure  silver,  preferably  foil,  place 
in  a  beaker,  add  lOc.c.  water  and  5c.c.  pure  nitric  acid,  and  cover  with 
clock  glass.  Warm  if  necessary,  and  after  all  is  dissolved  boil  gently. 
Wash  under  side  of  glass  and  the  inner  surface  of  beaker  till  beaker 
contains  about  75c.c.  liquid,  then  add  20c.c.  dilute  hydrochloric  acid 
and  stir.  Set  in  dark  place  till  ready  to  filter.  In  a  clean  Gooch  cru- 
cible make  a  mat  of  asbestos  with  aid  of  suction  of  the  filtering  pump 
as  shown  by  the  Instructor.  It  should  be  thick  enough  so  that  you 
cannot  see  light  through  it,  dry  for  half  an  hour  in  oven  at  about  140 
degrees,  cool,  in  desiccator  or  on  a  clean  surface  and  covered,  twenty 
minutes  and  weigh. 

With  suction  pump  filter  off  silver  chloride  into  crucible  as  shown, 
heat  in  oven  one  hour  at  about  140  degrees.  Desiccate  as  before  and 
weigh.  Find  weight  of  silver  chloride  and  thence  chlorine,  and  cal- 
culate the  equivalent  of  chlorine  if  the  equivalent  of  silver  is  107.9. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  25 

The  equivalents  of  bromine  and  iodine  may  be  determined  in  the 
same  way.  Indirectly  many  other  equivalents  may  be  determined  by 
the  method  as  the  following  problem  illustrates: 

If  3  grams  potassium  chloride  be  treated  with  an  excess  of  silver 
nitrate  and  the  silver  chloride  weighs  5.773  grams,  find  the  equivalent 
of  potassium,  if  that  of  chlorine  is  35.5. 

(b)  Weigh  accurately  a  small  porcelain  dish,  place  in  it  about 
0.5  gram  of  silver  foil  and  weigh  again.  Add  to  dish  5c.c.  con.  nitric 
acid  diluted  with  lOc.c.  of  water,  cover  with  a  watch  glass.  Give  it 
time  and  do  not  heat  unless  necessary.  When  all  the  metal  is  dis- 
solved, with  a  small  amount  of  water  in  a  fine  stream  from  the  wash 
bottle  wash  any  spattered  liquid  from  the  under  side  of  the  watch 
glass,  into  the  dish.  Add  to  the  dish  20c.c>  pure  dilute  HC1  and  eva- 
porate the  liquid  to  complete  dryness  on  a  water  bath.  Heat  the  dish 
in  an  oven  for  half  an  hour  at  about  125°,  or  heat  some  distance  above 
the  burner  flame  till  the  silver  chloride  begins  to  melt.  When  fully 
cold  weigh  accurately  the  dish  and  contents,  subtract  from  the  weight 
of  the  silver  chloride  the  weight  of  the  silver  and  find  the  equivalent 
cf  chlorine  calling  that  of  silver  107.9. 

42.— Equivalents  of  Copper  and  Other  Heavy  Metals :  The  equiva- 
lents of  copper  and  several  other  metals  may  be  determined  Toy  con- 
verting the  weighed  metal  into  nitrate  with  nitric  acid,  and  decom- 
posing the  nitrate  by  heat,  leaving  the  oxide. 

Use  accurately  weighed  copper  foil  or  clean  copper  turnings,  and 
proceed  the  same  as  in  41  (b),  but  use  no  hydrochloric  acid.  Evapor- 
ate to  complete  dryness,  and  heat  high  over  a  flame'  with  the  watch 
glass  on  the  dish.  Gradually  lower  the  dish  and  when  nearly  all  blue 
color  has  disappeared,  or  all  evidence  of  steam  in  the  case  of  other 
metals,  remove  the  watch  glass  and  apply  the  full  capacity  of  the 
burner  for  half  an  hour  when  there  should  remain  a  layer  of  black 
copper  oxide.  Material  spattered  upon  the  glass  must  be  washed  back 
into  the  cooled  dish  and  the  liquid  must  be  evaporated  again  on  the 
water  bath,  and  the  dish  strongly  heated.  Weigh  when  quite  cold. 

Prom  the  weight  of  the  copper  oxide  subtract  the  weight  of  the 
copper,  and  find  the  equivalent  of  copper  calling  oxygen  8.  If  16  is 
used  for  oxygen  the  number  obtained  for  copper  is  its  atomic  weight. 

MTROGEtf  A1VD  ITS  COMPOUNDS. 

43.— Preparation  of  Mtrogen  From  the  Air:  Fill  a  pneumatic 
trough  with  water  till  it  rises,  about  %  inch  over  the  shelf.  Place  up- 
on the  shelf  a  small  crucible  or  cupel  and  place  in  it  as  much  red 
phosphorus  as  will  lie  on  a  half  inch  of  the  end  of  a  knife  blade  or 
spatula.  Ignite  the  P  and  at  once  place  over  the  vessel  and  upon  the 
shelf  a  jar  or  wide  mouth  bottle.  Let  the  jar  remain  till  it  is  cool  and 


26 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


(he  white  fumes  have  been  absorbed.  Note  the  height  to  which  the 
water  has  risen  and  estimate  the  ratio  of  the  volume  of  oxygen  which 
has  been  consumed  to  the  nitrogen  which  remains.  Slip  cover  on  the 
jar,  place  it  upright  and  test  the  gas  remaining  with  a  burning  splint- 
er. What  other  gases  are  mixed  with  the  N  thus  obtained?  Why  does 
not  this  method  give  accurate  results  as  to  the  volume  of  oxygen  in 
air? 

4-1.— More  Accurate  Determination  of  Oxygen  in  Air:  Arrange  the 
apparatus  as  shown  in  fig.  11.  When  all  is  adjusted,  remove  ihe  test 
tube,  mix  about  7  cc.  of  a  solution  of  pyrogallol 
with  about  the  same  volume  of  sodium  hydroxide. 
At  once  pour  the  liquid  into  the  funnel.  By  open- 
ing the  clamp  let  the  liquid  down  till  it  just 
reaches  the  lower  end  of  the  tube.  Attach  the  test 
tube  by  inserting  the  stopper  firmly,  and  opening 
the  clamp  move  it  till  it  closes  out  of  the  way  upon 
the  rubber  and  glass  connecting  tube.  Move  the 
test  tube  to  a  horizontal  position  to  expose  the  air 
to  more  of  the  liquid  surface.  If  the  surface  of 
the  liquid  in  the  funnel  threatens  to  lower  into 
the  neck,  add  a  little  water.  When  you  are  sure 
the  liquid  has  ceased  to  enter  the  test  tube,  invert 
the  test  tube  and  bring  it  up  till  the  surfaces  of  the 
liquid  in  test  tube  and  funnel  are  at  the  same 
Fig.  11.  height,  as  shown  by  the  dotted  lines.  Estimate,  or 

better,  measure  the  lengths  of  the  columns  of  the  liquid  in  the  test 
tube  and  the  gas  which  remains,  Which  represents  oxygen  and  which 
nitrogen?  Find  their  ratio  by  volume  in  air. 

15. — Preparation  of  jVitrogen  from  Chemicals:  Set  up  the  appar- 
atus as  shown  in  fig.  12.  The  nitrogen  is  obtained  by  heating  a  solu- 
tion of  ammonium  nitrite  but  since  this  substance  is  difficult  to  make 
and  to  preserve,  the  same  results  may  be  obtained  by  heating  a  solu- 
tion of  ammonium  chloride,  NH4C1,  and  sodium  nitrite,  NaNO2.  These 
give  in  the  solution  the  same  groups,  NH4  and  NO2,  which  form  am- 
monium nitrite,  NH4NO2. 

Put  in  the  flask  10  grams  of  ammonium  chloride,  10  of  sodium  ni- 
trite and  lOOc.c.  of  water.  Heat  and  after  the  air  has  been  expelled  col- 
lect the  nitrogen  in  jars. 

Light  a  bit  of  candle,  place  it  upon  the  deflagration  spoon  and 
lower  it  into  a  jar  of  the  nitrogen.  Lower  burning  phosphorus  into 
another  jar. 

Note  the  curious  fact  in  this  experiment  that  the  oxygen  of  am- 
monium nitrite  oxidizes  the  hydrogen  of  the  same  compound;  that  is, 
one  part  cf  it  acts  as  a  reducing  agent  and  the  other  as  an  oxidizing 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


Fig.  12. 

agent.  Also,  hereafter  notice  that  whenever  oxidation  takes  place  witti 
respect  to  one  substance,  reduction  takes  place  at  the  same  time  with 
regard  to  some  other  substance.  Find  illustrations  of  this  under  prep- 
aration of  chlorine,  24. 

46. — An  analogous  method  may  be  used  for  the  preparation  of  N  by 
heating  an  intimate  mixture  of  ammonium  chloride,  2  grams,  and  po- 
tassium dichromate,  5  grams,  instead  of  ammonium  dichromate.  The 
method  of  heating  and  collecting  is  quite  the  same  as  that  for  oxygen 
in  6. 

4". — Preparation  of  Ammonia,  Place  about  a  half  gram  of  ammon- 
ium chloride,  ammonium  sulfate,  ammonium  nitrate  in  three  test 
tubes.  Add  to  each  about  2c.c.  of  sodium  hydroxide  solution  and 
warm.  Note  odor  in  each  case,  and  hold  in  each  tube  a  strip  of  moist 
turmeric  paper.  Repeat  using  about  the  same  amounts  of  the  am- 
monium compounds,  but  mix  each  with  about  its  own  weight  of  slaked 
lime.  Formulate  a  general  method  of  preparing  ammonia. 

Refer  to  a  text-book  for  an  account  of  calcium  cyanamide,  its 
manufacture,  hydrolysis,  use,  importance.  In  a  test  tube  place  about 
a  gram  of  the  substance,  barely  moisten  with  water.  Suspend  in  the 
tube  a  strip  of  moist  turmeric  paper,  corking  the  tube  and  thus  hold- 
ing the  strip  in  place.  Observe  evidence  of  ammonia  after  an  hour. 

Set  up  the  apparatus  as  in  fig.  13  and  provide  the  dry  fountain 
bottle  as  in  27.  The  collecting  jars  or  bottles  must  be  dry.  Weigh  and 


28 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


mix  in  mortar  10  grams  each  of 
ammonium  chloride  and  slaked 
lime.  Place  on  paper  and  slide 
into  iron  tube.  Protect  the  stop- 
per of  tube  with  wet  cotton. 
Heat  the  tube  and  collect  two 
bottles  or  jars  of  gas  by  the 
downward  displacement  of  air. 
When  full  moist  turmeric  paper 
held  at  the  mouth  of  a  jar  will 
be  instantly  turned  brown. 
While  it  is  bottom  upward 
clamp  cover  on  jar  or  cover 
bottle  with  glass  plate  ani 
place  on  desk  bottom  upward. 
Fill  the  fountain  bottle  and  pro- 
duce a  fountain  as  with  HC1  in 
Fig.  13.  27.  Turn  the  delivery  tube  down 

and  place  it  in  a  test  tube  containing  lOc.c.  pure  water  so  that  the  end 
of  the  delivery  tube  shall  reach  just  under  the  water.  Observe  ab- 
(sorption  of  the  ammonia  and  lower  the  delivery  tube  in  the  water  as 
necessary  to  absorb  all  the  gas.  Heat  as  long  as  ammonia  seems  to 
come  off,  remove  delivery  tube  from  test  tube  then  remove  the  lamp. 
Observe  odor  of  the  solution  of  ammonia  in  the  test  tube.  There 
is  evidently  gaseous  ammonia  above  the  liquid.  There  is  also  am- 
monia in  solution,  some  ammonium  hydroxide  and  some  ionized  am- 
monium hydroxide. 

48.— Action  of  Acids  and  Ammonia;  Neutralization:  In  a  jar  pour 
a  few  c.c  of  con.  HC1,  wet  the  sides  of  the  jar  and  pour  out  excess  oi7 
liquid.  Place  over  this  the  jar  of  ammonia  gas,  bottom  upward,  remove 
cover,  placing  jars  mouth  to  mouth.  Mix  by  reversing  the  pair,  placing 
the  jar  with  HC1  above.  The  fumes  and  the  solid  on  walls  of  jars  con- 
sist of  ammonium  chloride,  NH4C1,  made  by  direct  union  of  what?  Try 
a  burning  splinter  in  the  remaining  bottle  of  ammonia  gas. 

In  this  paragraph  and  in  all  other  such  cases  do  not  dip  test  papers 
into  liquids,  but  take  out  a  drop  of  the  liquid  with  a  stirring  rod  and 
touch  it  to  the  paper.  In  a  porcelain  dish  place  lOc.c.  of  ammonium 
hydroxide  and  neutralize  with  dilute  nitric  acid  in  just  the  same  way 
as  sodium  hydroxide  was  neutralized  with  HC1  in  28.  Evaporate  about 
two-thirds  of  the  liquid  over  a  flame  and  complete  the  evaporation  on 
a  water  bath.  Why  on  a  water  bath?  Press  dry  the  ammonium  nitrate 
between  folds  of  filter  paper. 

Small  quantities  of  ammonia  may  be  tested  for  as  directed  in 
Qualitative  Analysis,  Group  V.  Very  small  quantities  are  tested  for 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  29 

with  Nessler  solution,  which  is  a  strongly  alkaline  solution  of  mercury 
and  potassium  iodides.  If  at  hand  and  to  show  the  sensitiveness  of  the 
test,  fill  a  clean  test  tube  nearly  full  of  distilled  water,  and  in  another 
tube  of  distilled  water  dissolve  a  granule  of  ammonium  chloride  as 
large  as  a  pin  head.  To  each  tube  add  about  2c.c.  of  Nessler  solution. 
In  very  dilute  solutions  the  amounts  of  ammonia  are  in  proportion  to 
the  depths  of  color,  and  the  test  is  much  used  to  determine  very  small 
amounts  of  ammonia  in  water. 

49.— Preparation  of  Nitrogen  Monoxide  (Nitrous  Oxide):  Set  up 
the  apparatus  as  in  fig.  12,  but  omit  the  thistle  tube.  Heat  in  the  dry 
flask  15  grams  of  ammonium  nitrate,  and  regulate  the  heat  so  as  to 
control  the  flow  of  gas,  which  should  be  collected  in  three  jars  or  bot- 
tles. When  through  collecting  remove  the  delivery  tube  from  the 
water,  then  remove  the  burner. 

In  one  jar  thrust  a  glowing  splinter.  In  another  lower  burning 
phosforus.  Burn  a  piece  of  picture  cord  in  the  same  way  as  in  oxygen. 
Write  the  equation  and  compare' it  with  the  self-oxidation  of  ammonium 
nitrite  to  form  nitrogen. 

50. — Preparation  of  Xitrie  Acid:  Arrange  the  retort  and  receiving 
flask  as  in  the  preparation  of  bromine,  fig.  9,  but  retort  and  flask  should 
preferably  be  dry.  Put  into  the  retort  as  there  described  20  grams  of 
sodium  or  potassium  nitrate.  Add  through  the  tubulus  of  the  retort 
by  means  of  a  funnel  25c.c.  of  con.  sulfuric  acid.  At  once  wash  the 
funnel  and  measuring  cylinder.  Let  stand  till  acid  and  solid  are  in 
contact  throughout  and  then  heat  with  a  small  flame,  preferably  of  the 
crown  top.  There  is  no  need  of  cooling  the  receiving  flask  if  the  heat 
applied  is  properly  regulated.  Turn  down  the  flame  if  fumes  escape 
in  considerable  amount  from  the  flask.  Stop  heating  when  the  acid 
comes  over  slowly  and  the  liquid  in  the  retort  is  clear  and  seems  to 
be  viscid.  What  is  the  substance  left  in  the  retort?  To  remove  it  when 
cool  fill  the  retort  nearly  full  of  water  and  heat  gently,  best  on  a  water 
bath.  Pour  off  solution,  add  fresh  water  and  heat  again.  Repeat  till 
all  is  dissolved.  If  the  cake  of  solid  comes  loose  do  not  shake  it  about, 
since  the  thin  retort  will  be  broken. 

51.— Oxidizing  Action  of  Mtric  Acid:  Nitric  acid,  after  free  oxy- 
gen, is  the  most  important  of  all  oxidizing  agents  and  its  oxidizing  ac- 
tion should  be  well  understood. 

When  the  very  concentrated  acid  is  heated  a  part  of  it  breaks  down 
thus:  (1)  2HNO3=H2O+2NO2+O.  Whenever  the  more  dilute  acid  is  in 
contact  with  something  easily  oxidized  the  reaction  is  likely  to  be: 
(2)  2HNO3=H2O+2NO+3O.  To  illustrate,  (a)  heat  a  few  drops  of  your 
acid  and  note  the  red-brown  gas  given  off.  (b)  Place  a  loose  plug  of 
woolen  yarn  in  the  mouth  of  a  test  tube  which  contains  about  2c.c.  of 
your  acid,  and  boil  the  acid  till  the  vapors  ,set  the  wool  on  fire,  (c) 


30  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

Heat  a  little  sawdust  in  a  dish  till  it  begins  to  char  and  pour  upon  it 
about  Ic.c.  of  your  acid,  (d)  Place  about  3  grams  of  sugar  in  a  flask 
and  25  c.c.  of  con.  nitric  acid  from  the  shelf.  Heat  till  copious  fumes 
of  N02  are  formed.  The  sugar  is  mainly  oxidized  to  oxalic  acid.  The 
next  experiment  contains  a  good  illustration  of  oxidation  of  metals 
by'  HNO:i. 

In  the  most  common  case  of  oxidation  with  nitric  acid  the  acid 
itself  is  reduced  to  water  and  nitric  oxide,  NO.  Oxidation  and  reduc- 
tion, it  will  be  remembered,  go  on  together  usually.  In  some  cases  the 
reduction  does  not  go  so  far.  An  instance  is  the  reduction  of  a  nitrate 
to  nitrite  by  heating  with  lead  (see  54).  In  some  cases  the  reduction 
proceeds  to  the  formation  of  hydroxyl-amine,  HONH2,  but  more  often 
to  ammonia.  To  illustrate  this  treat  a  few  bits  of  aluminum  in  a  test 
tube  with  about  3  c.c.  of  sodium  hydroxide,  and  add  2  drops  of  con. 
nitric  acid.  Warm  the  tube  and  when  the  action  becomes  rapid  note 
odor  of  ammonia.  Incline  the  tube  and  hdld  in  it  moistened  turmeric 
paper  without  touching  the  glass.  Here  the  hydrogen  from  the  sodium 
hydroxide  and  Al  reduces  the  nitric  acid  to  ammonia.  Arrange  in  a 
horizontal  line  in  note  book  the  successive  reduction  products  of  ni- 
tric acid. 

Aqua  regia  is  a  mixture  of  1  part  of  con.  nitric  with  3  parts  of  con. 
hydrochloric  acid.  Make  about  1  cc.  of  the  mixture,  heat  it  and  notice 
chlorine  given  off.  It  will  dissolve  many  substances  that  are  attacked 
by  neither  of  the  acids  alone. 

52.— Action  of  Nitric  Acid  on  Metals,  Nitric  Oxide:  See  in  a  text 
book  the  electro-motive  series,  and  the  action  of  nitric  acid  on  metals. 

The  very  electro-positive  metals  will  give  some  hydrogen  with  di- 
lute nitric  acid.  Set  up  the  apparatus  to  prepare  H,  but  use  a  test 
tube  instead  of  the  flask.  Place  in  tube  magnesium  turnings.  Dilute 
con.  HN03  by  mixing  5  c.c.  with  30  c.c.  of  water.  Fill  two  test  tubes 
with  water  and  have  them  ready  to  collect  the  gas.  Put  in  the  dilute 
acid,  about  5  c.c.  at  a  time  and  after  the  air  has  been  expelled  collect 
two  test  tubes  of  the  gas.  Test  the  gas  for  H.  Note  red  gas  in  test 
tube  when  air  enters,  using  the  second  tube.  Save  the  liquid  in  the 
tube  in  which  H  was  evolved.  Repeat  the  experiment,  using  granulated 
zinc  instead  of  magnesium.  Can  you  detect  any  H?  Note  red  furnes  of 
nitrogen  dioxide,  N02,  always  formed  when  nitric  oxide,  NO,  comes  in 
contact  with  air  or  oxygen.  When  the  action  is  over  pour  off  a  little 
liquid  from  the  Zn  and  test  it  and  the  solution  saved  from  the  mag- 
nesium for  ammonia:  To  the  liquid  add  NaOH  in  excess,  hold  a  strip 
of  moist  turmeric  paper  in  the  tube  not  letting  it  touch  the  glass  and 
warm  gently.  Also  note  odor  of  ammonia. 

The  metals  lower  than  zinc  in  the  electro-motive  series  give  no  H 
with  nitric  acid,  and  those  below  hydrogen  do  not  give  H  with  any 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  31 

acid.  Try  the  action  of  con,  nitric  acid  diluted  with  an  equal  volume 
of  water,  on  a  little  tin,  copper,  powdered  antimony.  Oxides  are  ob- 
tained with  the  first  and  third  metals.  Would  HC1  act  on  the  copper 
and  antimony?  Could  HNO3  and  metals  be  used  to  prepare  pure  H? 

53.— Nitric  Oxide:  Use  apparatus  as  in  fig.  3.  Put  in  flask  25 
grams  granulated  copper,  about  10  cc.  water,  see  that  thistle  tube 
reaches  into  the  water,  then  add  con.  nitric  acid  as  necessary  to  secure 
a  moderate  flow  of  gas.  Reject  the  first  half  jar  full  of  gas.  Fill  three 
jars,  leaving  one  on  the  shelf  of  the  trough.  Pass  some  of  the  gas 
through  a  solution  of  ferrous  sulfate  in  a  tube.  The  dark  color  is  due 
to  FeSOiNO.  What  does  the  blue  solution  in  the  flask  contain?  Pour 
it  into  a  vessel  provided  for  the  purpose.  Wash  the  ^remaining  copper 
and  put  into  the  vessel  provided  for  it. 

Prepare  oxygen  as  in  6  or  9  and  pass  a  little  into  jar  on  shelf. 
The  red  gas  is  nitrogen  dioxide,  NO2.  Let  it  become  absorbed  then  pass 
in  more  oxygen.  Thus  continue  till  the  water  ceases  to  rise  in  the 
jar.  Write  the  volume  equation  for  the  union  of  NO  and  02.  What  is 
formed  when  NO2  dissolves  in  oold  water;  in  hot  water? 

In  one  jar  of  NO  burn  red  phosphorus.    Try  a  candle  in  the  other. 

Express  the  action  of  nitric  acid  on  copper  in  two  equations,  the 
first  of  which  shows  the  decomposition  of  the  acid  in  oxidation  as  in 
equation  (2),  51,  forming  3CuO;  the  second  the  dissolving  of  this  oxide 
in  the  acid. 

54.— Reduction  of  a  Nitrate  to  Nitrite,  Nitrous  Acid:  (a)  In  an 
iron  crucible  heat  and  stir  with  an  iron  rod,  spike  nail,  5  grams  of 
potassium  nitrate  and  20  grams  granulated  lead  till  the  mass  glows. 
When  cold  add  water  and  boil  for  some  time.  Filter  off  the  liquid  into 
a  tube.  Add  an  excess  of  dil.  H2SO4.  Test  the  gas  with  moist  starch 
iodide  paper.  The  sulfuric  acid  sets  free  nitrous  acid  which  breaks 
up  into  water,  NO  and  NO-;  equations.  Note  that  NO+NO2=N2O3,  the 
anhydride  of  nitrous  acid.  It  exists  only  as  a  liquid. 

(b)  Read  (a)  carefully.  In  a  flask  with  thistle  tube  and  right-an- 
gled delivery  tube  place  10 -grams  sodium  nitrite,  add  enough  water  to 
cover  end  of  thistle  tube,  place  delivery  tube  in  5  cc.  NaOH  in  a  wide 
test  tube.  Add  about  25  c.c.  dil.  sulfuric  acid  to  the  flask.  Nitrous  acid  is 
set  free,  but  breaks  up  as  in  (a) .  The  gases  are  absorbed  by  the  NaOH 
making  NaNO2.  Add  dilute  acid  to  this  solution  and  test  the  gas  as 
in  (a). 

55, — Nitrogen  Dioxide,  Nitrogen  Tetroxide:  The  lead  nitrate  used  in 
this  experiment  should  be  prepared  in  quantity  before  the  laboratory 
period  by  crushing  to  moderate  fineness  in  a  mortar  and  heating  for 
three  hours  in  an  oven  at  about  125°. 

Heat  10  grams  of  the  dried  lead  nitrate  in  iron  tube  shown  in  fig. 
13,  turning  the  delivery  tube  down,  and  placing  it  in  a  small  dry  flask. 


32  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

Heat  until  the  flask  is  filled  with  the  red-brown  gas,  NO2  mixed  with 
N2Oi.  Stopper  the  flask  and  heat  it  high  over  a  flame,  but  do  not  heat 
so  much  as  to  burst  the  flask.  A  temperature  of  142°  is  sufficient  to 
change  all  the  gas  to  NO2.  Note  depth  of  color.  Let  the  flask  cool  a 
few  moments  in  air,  then  under  running  water,  and  observe  the  depth 
of  color  again,  when  80%  of  the  gas  will  consist  of  colorless  N20u 
Write  the  equation  for  the  production  of  the  gas  and  the  reversible 
reaction  of  the  change  from  one  to  the  other  constituent  of  the  mixture. 

Heating  lead  nitrate  again  pass  the  gas  into  about  3  cc.  of  sodium 
hydroxide  for  a  few  moments,  then  add  dilute  sulfuric  acid  to  the 
solution.  What  evidence  have  you  that  a  nitrite  was  formed  by  the 
action  of  the  gas  on  sodium  hydroxide?  What  else  was  formed?  See 
text  book  for  the  action  of  this  gas  with  cold  water  and  with  hot  water. 

56.— Tests  for  Nitrogen  and  the  Nitrate  Radical:  (a)  Tests  for  ni- 
trogen in  organic  matter:  Heat  in  a  test  tube  with  soda-lime  or  a  mix- 
ture of  lime  and  powdered  sodium  hydroxide,  bits  of  woolen  material, 
such  as  yarn,  or  bits  of  leather,  dry  albumin  or  corn  meal,  and  hold  in 
the  mouth  of  the  tube  wet  turmeric  paper.  Is  ammonia  given  off? 

Secure  an  imperfect  test  tube  and  heat  in  it  any  of  the  organic 
substances  mentioned,  with  a  bit  of  sodium  not  larger  than  half  a  pea. 
(Do  not  touch  sodium  with  the  wet  hands).  Heat  strongly  till  the 
reaction  is  complete,  and  standing  well  back  place  the  hot  end  of  tube 
in  a  little  water  in  a  beaker.  The  end  cracks  off.  Boil  the  water  and 
filter.  To  a  portion  of  the  filtrate  add  a  little  solid  ferrous  sulfate, 
then  a  few  drops  of  ferric  chloride,  and  boil.  Now  make  acid  wiih 
dilute  HC1,  when  a  deep  blue  color  will  result.  To'the  rest  of  the  solu- 
tion add  a  crystal  of  potassium  nitro-ferri-cyanide,  which  will  give  a 
red  color,  showing  sulfur  in  the  organic  matter. 

(b)  Tests  for  the  nitrate  radical,  or  nitric  acid:  To  a  few  cc.  of  a 
solution  of  ferrous  sulfate  add  a  very  little  of  any  nitrate,  and  when 
the  substance  is  dissolved  and  the  solution  is  cool  incline  the  tube  and 
pour  in  steadily  con.  H2SO4,  so  that  it  will  run  down  and  collect  at  the 
bottom  of  the  tube.  Where  it  and  the  solution  meet  will  form  a  dark 
ring  of  FeSO.  (NO). 

The  following  test  should  be  used  only  for  excessively  small 
amounts  of  nitrate:  Upon  a  bulk  of  any  nitrate  not  larger  that  a  pin 
head  and  in  a  dish,  drop  about  ten  drops  of  phenol-disulphonic  acid. 
Warm  the  dish  gently,  best  on  a  water  bath  for  a  few  moments.  Add 
about  25  cc.  of  water  and  make  alkaline  with  ammonia.  A  yellow 
solution  shows  nitrate  radical  is  present. 

ELECTRO-CHEMISTRY—THEORY    OF    IONIZATION. 

57.— Electrical  Terms  and  Fundamental  Laws:  Using  good  text 
books  on  chemistry  and  on  physics  study,  till  clear  to  you,  the  meaning 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


33 


of  these  terms:     anode,  cathode,  Volt,  Ampere,  Ohm,  Coulomb,  Watt, 
electro-chemical  equivalent,  electro-motive  series  of  the  elements. 

What  is  Ohm's  Law?  On  a  110-Volt  circuit  find  the  current 
strength  or  amperage  transmitted  by  a  25-Watt  lamp.  Find  the  resis- 
ance  of  this  lamp  in  Ohms. 

What  is  Faraday's  Law?  What  are  the  electro-chemical  equiva- 
lents of  H,  O,  Cl,  SO4,  NO3,  Cu,  Ag,  Sn".  Sn""?  Give  relations  of 
electro-chemical  equivalents,  atomic  weights  (or  in  the  case  of  radicals 
the  sum  of  the  atomic  weights),  and  valences. 

•>*.— Electrical  Coiuluctiyity :  The  current  used  is  the  direct  light- 
ing circuit,  110  volts,  cut  down  by  a  25-Watt  lamp.  This  would  give 
what  amperage  through  a  wire  without  resistance  connecting  th,e  two 
binding  posts?  Calculate  and  read  the  ammeter.  Do  the  results  agree? 
Determine  the  comparative  conductivity  of  the  following  5th  normal 
solutions,  already  made  by  the  instructor:  hydrochloric  acid,  sulfuric 
acid,  acetic  acid,  sodium  acetate,  sodium  hydroxide  and  ammonium  hy- 
dioxicle.  Since  they  are  all  N/5  they  have  the  same  number  of  chemical 
equivalents  per  unit  volume.  According  to  whose  law  might  they  be 
expected  to  conduct  the  same? 

To  determine  the  conductivity  secure  one  of  the  conductivity  tubes 

as  shown  in  fig.  14,  wash,  rinse 
with  distilled  water  and  drain  it. 
In  each  case  be  careful  thus  to 
clean  the  tube.  Paste  a  strip 
of  gummed  paper  on  it  and 
in  each  case  fill  to  this 
mark.  Connect  the  tube  thus 
filled  with  the  binding  posts, 
note  the  intensity  of  the  glow  of 
the  lamp  and  read  the  ammeter. 
Thus  determine  the  conductivity 
of  all.  Why  should  they  conduct 
so  differently?  Knowing  the  re- 
sistance of  the  lamp  and  the  cur- 
rent it  alone  transmits,  calculate 
from  one  of  your  readings  of  the 
ammeter  the  resistance  of  the 
Fig.  14.  solution? 

Fill  the  conductivity  tube  with  a  neutral  solution  of  sodium  or  pot- 
assium sulfate,  add  a  few  drops  of  litmus  solution  and  mix  well,  and 
subject  the  solution  to  the  action  of  the  current.  Which  electrode 
gives  off  oxygen  and  shows  the  formation  of  acid  around  it;  which 
shows  hydrogen  evolved  and  alkali  formed  around  it?  Try  a  solution 
of  common  salt  without  litmus  solution.  Prove  with  starch-iodide  pa- 


34  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

per  that  Cl  is  given  off  at  the  anode.  Test  the  solution  at  the  cathode 
with  turmeric  paper.  By  what  secondary  reaction  is  alkali  formed  in 
the  last  two  instances?  What  is  the  action  of  sodium  on  water?  (13) 

In  electrolysis  many  of  the  heavy  metals  instead  of  acting  upon 
water  at  the  cathode  are  there  deposited.  Subject  a  solution  of  cop- 
per sulfate  to  electrolysis  for  several  minutes.  What  is  deposited  on 
the  cathode?  t  What  gas  is  given  off  at  the  anode?  Copper  sulfate 
itself  reacts  slightly  acid,  but  if  the  solution  at  the  anode  be  tested  it 
will  be  found  much  more  acid.  How  do  hydrogen  and  the  metals  ac- 
cumulate at  the  cathode  and  the  acid  radicals,  with  which  they  were 
associated,  at  the  anode? 

59. — Chemical  Facts  Best  Explained  on  the  Theory  of  lonization: 
The  Mature  of  Acids:  Test  with  blue  litmus  paper  dilute  HC1,  dilute 
HNO3,  dilute  H2S04,  acetic  acid  and  any  other  acids  that  may  be  avail- 
able. Why  should  substances  of  such  different  composition  all  turn 
the  paper  red?  Why  should  they  all  taste  sour?  Try  the  action  of 
each  on  bits  of  zinc.  Why  should  all  give  hydrogen?  Why  should 
they  all  neutralize  bases  giving  salts  and  water?  Why  should  they  all 
give  hydrogen  at  the  cathode  when  electrolyzed? 

It  would  seem  that  H  is  the  one  component  of  all  acids,  and  that 
to  it  their  acid  properties  are  due.  In  electrolysis  H  accumulates  at 
the  cathode  and  the  rest  of  the  acid  molecule  at  the  anode.  It  is  not 
far  to  the  thought  that  the  H  is,  in  the  solution  of  an  acid,  compara- 
tively free  from  the  rest  of  the  molecule.  Since  unlike  electric 
charges  attract  each  other  and  like  charges  repel,  the  inference  is  that 
the  H  atom  is  charged  positively  thus,  H+,  and  the  remainder  of  the 
molecule,  for  example,  Cl  is  charged  negatively,  thus,  Cl".  In  this  con- 
dition the  H  and  Cl  atoms  are  called  ions.  Ions  are  formed  when  acids, 
bases  and  salts  are  dissolved  in  water  and  without  regard  to  any  in- 
fluence of  the  electric  current. 

60.— The  Character  of  Bases:  Test  with  turmeric  paper  solutions 
of  sodium  hydroxide,  potassium  hydroxide,  ammonium  hydroxide,  cal- 
cium hydroxide,  barium  hydroxide.  Again,  why  should  such  different 
substances  all  turn  the  paper  brown?  They  also  all  neutralize  acids  in 
the  same  way  forming  salts  and  water.  Their  common  constituent  is 
hydroxyl  OH.  Since  all  bases  have  OH  it  is  inferred  that  the  alka- 
linity is  due  to  this  group.  It  seems  to  be  easily  separable  from  the 
metal  of  such  group  as  NH4.  In  electrolysis  it  travels  to  the  anode 
where  it  breaks  up  into  water  and  oxygen  while  the  metal  goes  to  the 
cathode.  For  reasons  given  under  acids,  it  is  believed  that  a  base  in 
solution  is  more  or  less  ionized.  That  is,  sodium  hydroxide  consists 
largely  of  the  ions  Na+  and  OH". 

61. — Ready  Cleavage  or  lonization  of  Salts:  In  as  many  test  tubes 
place  a  few  drops  of  solutions  of  the  following:  NaCl,  KC1,  NH4C1,  CaCla 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  35 

BaCl2  and  other  chlorides  that  may  be  at  hand.  Add  to  each  a  few 
drops  of  silver  nitrate.  Why  should  all  these  different  chlorides  gives 
the  same  precipitate  of  silver  chloride?  Silver  sulfate  or  silver  ace- 
tate might  have  been  used  instead  of  silver  nitrate  and  precisely  the 
same  precipitate  of  silver  chloride  would  have  resulted.  In  the  solu- 
tions of  chlorides  it  is  evident  that  chlorine  is  very  slightly  held  by  the 
metals  if  at  all,  and  the  same  is  true  of  the  silver  in  the  solutions  of  the 
silver  salts.  All  these  salts  are  electrolytes,  the  metals  going  to  the  ca- 
thode and  the  non-metals  to  the  anode.  The  inference  is  that  they  are 
largely  ionized  in  solution.  Thus,  sodium  chloride  is  largely  Na+  and 
d~,  and  silver  nitrate  is  Ag+  and  NO3~. 

Any  element  as  Cl  is  not  always  an  ion  in  solution.  To  solutions  of 
pure  potassium  chlorate,  chloral  and  chloracetic  acid  add  a  little  silver 
nitrate  solution.  No  silver  chloride  is  obtained,  since  KC1O3  ionizes  into 
K+  and  C103",  chloracetic  acid  into  H+  and  CH2C1CO2~,  while  chloral 
hydrate  gives  no  ions. 

62. — Degree  of  lonization:  Refer  to  your  experiment  on  conduc- 
tivity (58).  Did  all  the  acids  conduct  equally  well  and  were  they 
equally  ionized?  Compare  the  conductivity  and  ionization  of  sodium 
hydroxide  and  ammonium  hydroxide ;  of  acetic  acid  and  sodyAm  acetate. 

The  ionization  of  the  same  substance  may  be  increased  by  diluting 
and  decreased  by  concentrating  its  solution. 

(a)  Compare  the  colors  of  2-normal  solutions  of  copper  sulfate, 
nitrate  and  chloride.     The  color  of  each  is  supposed  to  be  due  to  the 
color  of  the  undisscciated  salt  and  to  the  copper  ion.     Now  dilute  a 
small  portion  of  each  solution  with  10  times  its  volume  of  water.    Why 
are  the  solutions  now  more  nearly  the  same  color?    To  a  portion  of 
each  solution  add  an  excess  of  ammonia.    The  same  color  is  due  to  the 
same  ion,  Cu(NH3)4++, 

(b)  The  concentration  of  a  solution  may  be  in  effect  increased 
and  the  dissociation  of  the  solute  decreased  by  adding  a  substance  hav- 
ing an  ion  in  common  with  the  solute.     Solid  copper  bromide  is  black, 
its  concentrated  solution  is  brown  due  to  CuBr2,  while  its  dilute  solu- 
tion as  that  of  every  other  cupric  salt  shows  blue  due  to  the  Cu  ion.  In 
a  dish  dilute  about  2  c.c.  of  the  brown  solution  till  it  becomes  blue.  To 
one  half  of  the  blue  solution  add  a  few  drops  of  con.  HBr,  which  has 
the  common  ion  Br,  till  brown.       To  the  other  half  add  solid  copper 
chloride,  having  the  common  ion  Cu,  till  brown.    Explain  how  the  ad- 
dition of  the  common  ion  effects  these  changes. 

(c)  Dilute  a  little  acetic  acid  from  shelf  bottle  with  20  times  its 
volume  of  water.     Add  a  few  drops  of  methyl  orange  and  divide  into 
two  portions.     To  one  add  one  half  its  volume  of  con.  sodium  acetate 
and  compare  its  color  with  that  of  the  other  half.    What  common  ion 
was  added  and  what  was  the  effect  on  the  acidity  of  the  acid? 


36  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

(d)  To  a  saturated  solution  of  salt  add  concentrated  HC1  till  a 
large  precipitate  of  salt  is  obtained,  and  account  for  its  formation. 

63. — Acidity  Due  to  Hydrolysis:  Many  substances  containing  no 
hydrogen  give  an  acid  solution.  They  are  mostly  salts  of  metals  which 
form  weak  bases  and  radicals  of  stronger  acids.  With  blue  litmus  pa- 
per test  solutions  of  salts  of  copper,  iron,  aluminium.  In  these  cases 
the  hydrolysis  goes  only  a  little  way,  for  example  thus: 

FeCl3+H20=(reversibly)FeCl,OH+HCl. 

Add  a  little  antimony  chloride  and  bismuth  chloride  to  five  times 
their  volumes  of  water.  Here  the  reactions  go  far  to  the  right  precipi- 
tating SbOCl  and  BiOCl.  Now  reverse  the  reactions  by  adding  con. 
HC1.  See  the  hydrolysis  of  the  halides  of  phosphorous,  34  and  79. 

64. — Alkalinity  Due  to  Hydrolysis:  Some  salts  of  strongly  basic 
metals  and  weak  acids  give  alkaline  solutions  due  to  hydrolysis.  Test 
with  turmeric  paper  or  a  few  drops  of  phenol-phthalein  solutions  of  bor- 
ax, sodium  carbonate,  Na2CO3,  and  sodium  phosphate,  HNa2POi.  In  the 
case  of  sodium  carbonate,  Na2CO3+H2O=HNaC034-NaOH,  and  of 
course  the  last  is  highly  ionized,  and  the  alkalinity  is  due  to  the  ion 
OH~.  For  olher  examples  see  77  and  95. 

65. — The  Electro-Motive  Series  of  Elements:  In  the  following  any 
metal  which  replaces  another  from  solution  is  said  to  have  a  greater 
solution  tension  or  to  be  more  electro-positive.  In  a  little  zinc  sulfare 
solution  place  bits  of  magnesium  turnings  and  let  stand.  When  the 
action  has  nearly  ceased  note  gray  deposit  of  zinc.  Place  zinc  in  a  so- 
lution of  cadmium  chloride  and  later  note  cadmium  deposited  on  the 
remaining  zinc.  Try  cadmium  or  zinc  in  solutions  of  copper  sulfate 
and  lead  acetate  and  note  copper  and  lead  deposited.  Try  iron  in  a 
copper  solution.  Place  a  strip  of  copper  in  a  solution  of  mercuric 
chloride  and  after  a  few  minutes  remove,  rub  and  note  mercury  coat- 
ing. Put  a  copper  wire  in  a  silver  solution  and  note  deposit  of  silver, 
Try  silver  in  a  solution  of  gold  chloride  and  note  deposit  of  gold  on 
the  silver.  All  these  cases  are  practically  alike.  One  metal  goes  into 
solution  as  ions  and  the  other  is  as  it  were  forced  out  as  neutral  metal; 
for  example,  Cu+++SO4-  +  Zn=(reversibly)Zn+++SO4"+Cu.  Write  a 
similar  ionic  equation  for  the  preparation  of  H  with  zinc  and  sulfuric 
acid.  Has  the  SO4  much  to  do  with  either  case? 

Arrange  the  above  metals  in  the  order  of  their  capacities  to  replace 
other  metals,  or  their  ionizing  tendency.  Compare  the  result  with  the 
arrangement  of  the  Electro-Motive  Series  in  a  text-book.  Wh'at  two 
metals  above  if  used  for  the  plates  in  an  electric  battery  cell  would 
give  the  greatest  electromotive  force? 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  37 

SULFUK. 

<;#.— Solubility  and  Crystallization:  Do  not  heat  carbon  disulfide 
or  have  a  flame  near  it,  since  it  is  very  volatile  and  easily  inflam- 
mable. 

Place  in  a  dry  test  tube  about  2  grams  of  powdered  sulfur,  add 
not  more  than  one-fourth  test  tube  full  of  carbon  disulfide,  shake  and 
filter  through  a  dry  filter  into  a  crystallizing  dish  or  small  beaker. 
Let  two  or  three  drops  fall  into  a  clean  watch  glass.  After  the  liquid^ 
has  evaporated  at  room  temperature,  examine  the  crystals  in  each  ves- 
sel, and  those  in  the  watch  glass  with  a  microscope.  The  crystals  be- 
long to  the  rhombic  variety  of  sulfur. 

<J7. — Effects  of  Heat,  Allotropic  Forms:  Since  the  test  tube  can- 
not be  cleaned  fill  an  imperfect  one  three-fourths  full  of  lumps  of 
sulfur,  slowly  and  evenly  heat  till  it  is  all  changed  to  a  yellow,  mobile 
liquid.  Now  increase  the  temperature  and  note  that  it  becomes  darker. 
Find  a  temperature  at  which  the  tube  may  be  inverted  for  a  moment 
without  the  sulfur's  running  out  Increase  the  temperature  till  it 
shows  signs  of  boiling  and  becomes  quite  fluid  again.  Pour  in  a  steady 
small  stream  into  a  beaker  or  other 'vessel  full  of  water.  Examine 
this  "plastic  sulfur",  place  aside  and  examine  at  the  end  of  the  lab- 
oratory period  and  at  the  next  period.  It  gradually  changes  to  mono- 
clinic  sulfur,  and  after  a  long  time  to  rhombic,  the  permanent  form. 
Consult  text-book  on  the  forms  of  sulfur,  and  their  properties. 

<js. — Hydrogen  Sulfide:  Fit  a  small  bottle  or  flask,  the  former  pre- 
ferred, with  a  delivery  tube  consisting  of  two  right  angled  pieces  joined 
with  rubber  tubing,  the  outer  piece  long  enough  to  reach  to  the  bot- 
tom of  the  collecting  vessel  or  test  tube  when  resting,  on  the  desk. 
Place  in  the  generating  vessel  20  grams  of  ferrous  sulfide  in  small 
bits,  sliding  in  the  larger  pieces  to  avoid  breakage  if  a  flask  is  used, 
add  about  40  c.c.  of  water  and  drop  directly  upon  it  a  little  at  a  time 
con.  sulfuric  till  a  suitable  flow  of  gas  is  secured.  If  the  flow  of  gas 
becomes  too  slow  at  any  time  do  not  add  more  acid,  but  pour  off  the 
liquid  which  is  nearly  exhausted  of  acid  but  contains  much  iron  salt, 
and  add  water  and  acid  as  in  the  beginning. 

Collect  a  bottle  of  the  gas  by  displacement  of  air,  in  the  same 
way  as  chlorine,  and  burn  the  gas  in  the  bottle.  What  is  the  deposit 
on  the  sides  of  the  bottle?  Collect  gas  in  a  test-tube  one- fourth  full 
of  water,  cover  mouth  of  tube  with  the  thumb  and  shake.  Is  the  gas 
soluble  in  water?  Turn  upward  the  outer  section  of  the  delivery  tube 
and  light  the  gas  at  the  orifice.  Hold  cold  porcelain  #s  a  dish  or  cru- 
cible lid  in  the  flame  at  about  its  middle.  What  is  the  yellow  deposit? 
Insert  the  flame  into  a  bottle  with  wide  mouth  as  far  as  you  can  with- 


38  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

out  extinguishing  the  flame.  Why  does  sulfur  deposit  on  the  dish  and 
bottle  instead  of  burning? 

69.— Precipitation  with  Hydrogen  Sulfide:  Hydrogen  sulfide  is 
much  used  in  separating  the  metals  into  groups  in  analytical  chemis- 
try. In  test  tubes  in  the  test  tube  holder  place  about  5  c.c.  of  solu- 
tions of  salts  of  the  following  metals,  using  chlorides,  nitrates,  ace- 
tates as  may  be  convenient:  silver,  lead,  mercury,  bismuth,  copper, 
cadmium,  antimony,  arsenic,  tin;  also,  zinc,  cobalt,  nickel,  manganese, 
iron;  also,  barium,  calcium,  strontium,  sodium,  add  a  drop  or  two  of 
dilute  HC1  or  HNOs  to  each.  Pass  the  gas  flowing  with  moderate  ra- 
pidity into  each  of  these  solutions  for  about  a  minute,  washing  off  the 
delivery  tube  by  dipping  into  water  before  inserting  it  into  a  new  so- 
lution. Note  carefully  what  metals  can  be  precipitated  as  sulfides 
with  H2S  from  slightly  acid  solutions  remembering  the  acid  added  and 
the  fact  that  when  precipitation  occurs  acid  is  necessarily  set  free. 
The  precipitation  of  zinc  is  not  complete.  Pass  an  excess  of  gas  into 
the  tube  containing  zinc,  filter  off  the  precipitate  and  add  to  the  fil- 
trate and  to  each  succeeding  solution,  beginning  with  cobalt,  ammon- 
ium hydroxide,  which  should  give  precipitates  of  sulfides  in  all  until 
barium  is  reached.  That  is,  several  metals  not  precipitated  as  sul- 
fides in  acid  solution  are  precipitated  with  hydrogen  sulfide  and  an 
alkali  and  they  form  another  group.  Name  the  metals  not  precipi- 
tated as  sulfides. 

Return  to  the  sulfides  precipitated  from  acid  solution,  pour  off  the 
liquid  from  the  sulfides  of  Hg,  Cu,-  Bi  and  try  to  dissolve  each  by  boil- 
ing'with  dilute  nitric  acid.  How  may  Hg  be  separated  from  the  other 
two?  Pour  off  the  liquid  from  the  sulfides  of  lead,  antimony  and  ar- 
senic and  heat  each  with  about  5  c.c.  of  ammonium  sulfide,  but  do  not 
boil.  How  may  lead  sulfide  be  separated  from  arsenic  and  antimony 
sulfides? 

These  are  merely  illustrations  of  the  use  of  hydrogen  sulfide  and 
many  others  will  be  met.  Instead  of  using  hydrogen  sulfide  and  am- 
monium hydroxide  separately  to  precipitate  certain  sulfides,  it  is  com- 
mon practice  to  use  ammonium  sulfide  which  may  be  easily  made  by 
saturating  diluted  ammonium  hydroxide  solution  with  hydrogen  sul- 
fide, and  then  adding  an  equal  volume  of  the  same  ammonia  solution 
which  gives  essentially  (NH4)2S. 

70. — Reducing  Action  of  Hydrogen  Snlfide.  Hydrogen  sulfide  is 
not  a  very  stable  compound  and  it  acts  as  a  reducing  agent  in  much 
the  same  way  as  hydriodic  acid.  Pass  hydrogen  sulfide  into  a  dilute 
acidified  solution  of  potassium  dichromate,  till  green ;  into  a  dilute  acid- 
ified solution  of  potassium  permanganate  till  colorless;  into  bromine 
water  and  a  dilute  solution  of  iodine  till  colorless.  In  the  last  two 
cases  filter  off  the  sulfur  and  test  liquids  with  litmus  paper.  What 


LABORATORY  MANUAL,  OF  GENERAL  CHEMISTRY 


39 


acids  were  formed?  Compare  the  reducing  action  of  hydrogen  sulfide 
with  that  of  sulfur  dioxide  in  experiment  71. 

71. — Sulfur  Dioxide:  (Read  reference  book  on  sulfur  dioxide,  sul- 
fur trioxide  and  sulfuric  acid.)  In  a  test  tube  treat  a  little  sodium 
sulfite  with  dilute  acid.  Note  odor  and  test  with  a  strip  of  paper  mois- 
tened with  mercurous  nitrate.  It  should  turn  dark. 

Determine  whether  bits  of  charcoal  and  sulfur  will  give  SO2  when 
heated  with  con.  sulfuric  acid.  Explain  their  action  on  the  acid.  Ar- 
range apparatus  as  in  the  preparation  of  hydrochloric  acid.  Place  in 
the  flask  20  grams  of  granulated  copper  or  bits  of  sheet  copper  or 
turnings  and  add  25  c.c.  con.  sulfuric  acid.  Before  heating  read 
through  the  experiment  and  have  everything  needed  at  hand  so  as  to 
continue  the  flow  of  gas  the  minimum  of  time. 

Heat  till  gas  comes  off  freely  and  control  the  rate  by  turning 
down  the  flame.  After  acidifying  them  with  acetic  acid  pass  gas  into 
a  solution  of  potassium  dichromate  till  green  and  into  a  solution  of 
potassium  permanganate  till  colorless;  also,  into  bromine  water  till 
colorless,  into  dilute  nitric  acid,  and  into  distilled  water.  To  each  so- 
lution add  barium  nitrate  and  then  dilute  HC1.  The  formation  of  a 
white  precipitate  by  barium  nitrate,  insoluble  in  HC1  shows  the  pres- 
ence of  the  sulfate  radical,  SO4. 

The  first  action  of  SO2  on  passing  it  into  the  above  solutions  was 
probably  the  formation  of  the  unstable  sulfurous  acid,  H2SO3.  All  the 
substances  into  which  the  gas  was  passed  were  oxidizing  agents  except 
of  course  the  distilled  water.  Show  their  action  on  the  sulfurous  acid. 
Test  for  the  sulfate  radical  in  the  dilute  sulfuric  acid  on  the  shelf,  in 
solutions  of  magnesium  sulfate,  copper  sulfate  and  other  sulfates  that 
may  be  at  hand. 

72.— Preparation  of  Sulfuric  Acid:  Set  up  the  apparatus  as  shown 
in  fig.  15.  The  ignition  tube  may  be  of  hard  glass  or  of  half  inch  gas 
pipe.  If  the  latter  is  used  it  should  be  12  to  14  inches  long  to  prevent 
the  overheating  of  the  stoppers.  Even  then  it  is  well  to  spray  the  ends 
with  the  wash  bottle  catching  the  water  as  it  runs  off,  or  to  wrap  the 
ends  with  cotton  or  filter  paper  and  keep  it  saturated  with  water. 


Fig.  15. 


40  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

Using  loose  plugs  of  asbestos  fiber  or  steel  wool  ^confine  near  the 
middle  of  the  tube  about  15  grams  of  granular  iron  pyrite.  Connect 
the  outer  end  of  the  tube  with  the  compressed  air  system.  If  this  is 
not  available  a  bulb  pump,  foot-bellows  or  even  a  bicycle  pump  may  be 
used  to  force  through  the  air.  Instead  of  forcing  through  the  air  the 
other  end  of  the  apparatus  may  be  connected  with  the  exhaust 
system  or  a  filtering  pump. 

Place  in  the  test  tube  3  c.c.  of  con.  nitric  acid,  preferably  fuming 
acid,  and  to  the  larger  bottle,  connected  with  it  add  about  5  c.c.  of 
water.  The  smaller  bottle  contains  a  solution  of  sodium  hydroxide  to 
absorb  the  excess  of  noxious  gases, 

Heat  the  tube  containing  the  pyrite,  preferably  with  a  wing  burner 
and  force  or  draw  through  a  moderate  stream  of  air.  When  the  bot- 
tles are  filled  with  white  fumes  it  may  be  known  that  the  pyrite  is  oxi- 
dizing. Moderate  the  heat  to  avoid  distilling  off  unburned  sulfur,  and 
regulate  the  air  so  that  only  a  small  amount  of  fumes  escape  from  the 
smaller  bottle.  Continue  the  operation  10  to  20  minutes.  To  avoid 
sucking  back  remove  the  air  connection  and  then  the  burner. 

Transfer  the  liquids  in  the  test  tube  and  larger  bottle  to  a 
weighed  porcelain  dish,  rinse  with  a  little  water  and  add  this  to  the 
dish.  With  a  free  flame  evaporate  till  the  steam  and  nitric  acid  va- 
por give  place  to  the  dense  gray  fumes  of  sulfuric  acid.  Let  the  dish 
cool  and  weigh  it.  Three  to  four  grams  of  con.  acid  should  be  made  in 
a  10  minutes  run.  Show  that  the  acid  will  carbonize  a  splinter  dipped 
into  it.  Dilute  a  few  drops  of  it  and  test  with  barium  chloride.  Re- 
move and  examine  the  contents  of  the  ignition  tube.  What  results 
from  burning  iron  pyrite  in  air? 

73.— Sodium  Sullite  ami  Thiosulfate:  Dissolve  10  grams  anhydrous 
sodium  carbonate  in  50  c.c.  of  water  with  aid  of  heat  if  necessary,  cool 
and  place  just  one  half  of  the  volume  in  a  large  test  tube.  Pass  into 
the  solution  in  tube  sulfur  dioxide  made  by  heating  in  a  flask  with 
thistle  tube  25  grams  of  copper  turnings  with  25  cc.  of  concentrated 
sulfuric  acid.  The  delivery  tube  should  reach  to  the  bottom  of  the 
test  tube.  When  the  flow  of  gas  begins  the  heat  should  be  reduced  so 
as  to  avoid  a  too  rapid  stream.  As  the  carbon  dioxide  is  set  free  the 
liquid  shows  a  tendency  to  froth.  The  reaction  is  completed  in  about 
15  minutes  when  small  bubbles  disappear,  the  sulfur  dioxide  appar- 
ently is  no  longer  absorbed,  and  the  liquid  smells  strongly  of  the  gas. 
The  result  is  acid  sodium  sulfite.  Now  add  this  solution  to  the  re- 
served half  of  the  sodium  carbonate  solution  which  will  give  normal 
sodium  sulfite.  Why?  Place  the  combined  solutions  in  a  dish  and  boil 
for  two  or  three  minutes.  If  evaporated  nearly  to  dryness  on  a  water 
bath  and  allowed  to  cool  crystals  of  sodium  sulfite  would  be  obtained. 
The  whole  of  the  salt  may,  however,  be  converted  into  sodium  thiosul- 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  41 

fate.  For  .this  purpose  add  to  the  solution  three  grams  of  powdered 
sulfur  and  boil  gently  for  twenty  minutes  replacing  from  time  to  time 
the  water  evaporated.  Filter  off  the  sulfur  and  evaporate  the  filtrate 
over  a  water  bath  till  crystals  begin  to  form  and  let  the  solution  be- 
come cold. 

Pour  off  the  solution  from  the  crystals  and  use  it  to  study  the 
properties  of  sodium  thiosulfate. 

To  a  small  volume  of  a  solution  of  iodine  add  the  solution  of  thio- 
sulfate till  the  iodine  solution  is  decolorized. 

Make  a  little  silver  chloride  by  adding  a  few  drops  of  a  solution 
of  some  chloride  to  a  few  drops  of  silver  nitrate.  Now  add  the  thio- 
sulfate solution,  a  little  at  a  time  with  shaking,  till  the  silver  chloride 
dissolves.  What  is  the  use  of  sodium  thiosulfate,  "hypo,"  in  photo- 
graphy ? 

To  the  remainder  of  the  solution  of  thiosulfate  add  an  excess  of 
dilute  hydrochloric  acid,  warm  and  note  the  precipitation  of  sulfur 
and  the  odor  of  sulfur  dioxide.  Account  for  the  gas  and  the  sulfur. 

Here  as  elsewhere  study  the  experiment  with  aid  of  a  reference 
book  till  all  is  clear  and  write  the  equations. 

PHOSPHORUS. 

70:  ..White  and  Red  Phosphorus,  Allotropic  Forms:  (Caution: 
Handle  white  phosphorus  only  with  the  forceps.  Keep  trace  of  all 
pieces  used  and  any  remaining  must  not  be  left  about  the  laboratory 
desk  or  thrown  in  waste  jars,  but  put  into  a  vessel  of  water  provided 
for  that  purpose). 

Place  a  piece  of  white  phosphorus  as  large  as  a*grain  of  wheat  in 
water  in  a  test  tube  and  boil.  Phosphorus  vapor  goes  off  with  the 
water  vapor  but  burns  Jo  oxide  on  reaching  the  air.  Try  red  phos- 
phorus in  the  same  way.  Does  it  volatilize? 

Heat  a  little  red  P  in  a  dry  narrow  test  tube  and  note  sublimate 
of  white  P  on  wall  of  tube.  Change  white  P  to  red  P  by  dropping  a 
small  crystal  of  iodine  upon  white  P  in  a  dry  test  tube  and  warming. 

Dissolve  a  little  white  P  in  about  3  c.c.  of  carbon  disulfide  in  a  dry 
test  tube,  pour  the  liquid  upon  filter  paper,  place  this  upon  the  ring 
of  a  retort  stand  and  let  dry.  Try  to  dissolve  red  P  in  a  little  carbon 
disulfide.  Do  not  have  a  flame  near  carbon  disulfide.  What  is  meant 
by  "allotropic  forms"? 

"7.— Phosphine:  Drop  a  piece  of  fresh  calcium  phosphide  into  a 
small  bottle  or  beaker  full  of  water.  The  gas  set  free  is  mainly  phos- 
phine,  PH3,  but  it  is  mixed  with  another  hydride,  P2H4,  which  takes  fire 
spontaneously. 

78. —  (Caution:  The  following  experiment  is  a  beautiful  one,  but 
it  is  attended  with  some  danger.  It  may  be  undertaken  by  a  small 


42  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

class  under  constant  supervision) .  Set  up  apparatus  as  in  the  prepara- 
tion of  nitrogen,  but  place  the  end  of  delivery  tube  in  a  dish  of  water, 
fig.  12.  Place  in  the  flask  about  20  cc.  of  sodium  hydroxide,  made  by 
dissolving  the  solid  in  three  times  its  weight  of  water.  Add  two  or 
three  small  bits  of  white  P,  then  about  1  c.c.  of  ether  having  no  flame 
near,  stopper  the  flask  and  heat.  When  the  gas  begins  to  ignite  on 
reaching  the  air,  moderate  the  heat  so  as  to  maintain  a  rate  of  one 
bubble  every  two  or  three  seconds.  Carefully  insert  the  delivery  tube 
into  a  test  tube  of  alcohol  by  which  the  P2H4  is  dissolved.  Does  the  es- 
caping gas  now  take  fire? 

When  cold  carefully  remove  the  stopper,  fill  flask  with  water,  wash 
the  remaining  P  and  place  in  vessel  of  water  provided. 

79.— Chlorides  of  Phosphorus  and  Hydrolysis.  Refer  to  34,  63.  In 
a  dish  place  one  drop  of  water  and  add  one  drop  of  phosphorus  tri- 
chloride, best  from  a  dropping  bottle.  What  acid  gas  is  given  off? 
What  acid  of  P  remains?  In  another  dish  place  about  the  same  amount 
of  water  and  phosphorus  pentachloride  of  about  the  bulk  of  a  pea, 
handling  it  with  the  spatula.  What  gas  is  here  given  off?  Save  the 
liquid  to  test  for  phosphoric  acid  in  the  next  experiment. 

80.— Tests  for  the  Radical  P04  and  Phosphoric  Acids:     (a)     To  a 

few  c.c.  of  sodium  phosphate  add  one-fourth  its  volume  of  ammonium 
chloride,  make  alkaline  with  ammonia  and  then  add  magnesium  sul- 
fate.  The  slowly  forming  precipitate  is  ammonium  magnesium  phos- 
phate, (NH4)  Mg  P04,  and  this  is  a  good  test  for  phosphoric  acid  pro- 
vided arsenic  acid  is  not  present.  Add  water  to  the  second  dish  in  the 
preceding  experiment,  boil,  transfer  to  a  test  tube  and  test  for  PO4. 

(b)  To  test  for  phosphate  in  a  substance  insoluble  in  water  dis- 
solve  a  small  portion  in  dilute  nitric  acid,  and  to  a  few  drops  of  the 
solution  add  twice  its  volume  of  ammonium  molybdate  and  warm.    A 
yellow  precipitate  shows  phosphate,  provided  arsenic  acid  is  known  to 
be  absent.    If  practicable  secure  a  piece  of  bone,  place  it  in  dilute  ni- 
tric acid,  and  near  the  end  of  the  period  test  the  solution  for  phos- 
phate by  this  method. 

(c)  To  solutions  of  sodium  phosphate  HNa2  PO4  and  "microcosmic 
salt,   NH4HNaP04,  add   silver  nitrate.     Note  color  of  precipitate  the 
same  in  both  cases  and  try  its  solubility  by  adding  to  one  tube  dilute 
nitric  acid  and  to  the  other  an  excess  of  ammonia. 

81. — Other  Acids  of  Phosphorus:  In  a  crucible  heat  a  small 
amount  of  ordinary  solid  sodium  phosphate  and  after  the  water  of 
crystallization  has  been  driven  off  continue  with  full  flame  for  five 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  43 

minutes.  When  cold  dissolve  some  of  the  fused  mass  and  test  with 
silver  nitrate.  The  ordinary  sodium  phosphate,  Na2HPO4  has  been 
changed  to  sodium  pyrophosphate,  NaiPaOi.  Repeat  this  experi- 
ment using  microcosmic  salt.  While  the  water  is  evaporating 
test  for  ammonia  with  turmeric  paper,  and  note  odor.  This  salt  is 
changed  by  heating  to  sodium  metaphosphate,  NaPOs,  at  high  tempera- 
ture. Let  the  crucible  cool,  dissolve  most  of  the  metaphosphate  and 
test  with  silver  nitrate.  Compare  the  silver  tests  for  phosphate,  pyro- 
phosphate and  meta-phosphate. 

ARSENIC. 

S2. — Acidify  strongly  with  dilute  HC1,  dilute  solutions  of  sodium 
arsenite  NasAsOj  and  arsenate  NasAsO4  and  without  heating  pass  into 
each  H2S,  for  a  few  moments.  Do  you  get  a  precipitate  in  the  second 
solution?  Now  heat  each  nearly  to  boiling  and  pass  in  the  gas  again 
till  the  precipitation  is  complete.  Filter  on3  the  arsenic  sulfide,  As2S^, 
and  wash  it  on  the  filter.  Try  to  dissolve  a  part  of  the  substance  in 
con.  HC1  in  a  dish.  Now,  boiling  only  gently  add  very  small  bits  of 
potassium  chlorate  from  time  to  time  till  nearly  all  the1  sulfide  is  dis- 
solved. Add  a  little  more  con.  HC1  to  replace  that  boiled  away,  if  ne- 
cessary. Make  alkaline  with  ammonia  the  solution  in  the  dish  and  fil- 
ter, and  to  the  filtrate  add  a  solution  of  magnesium  sulfate  or  chlo- 
ride. Compare  this  test  for  arsenate  with  that  in  the  next  section,  and 
with  the  test  for  phosphoric  acid.  (80a.) 

Remove  most  of  the  remaining  sulfide  to  a  clean  dish,  add  about  o 
c.c.  of  ammonium  sulfide  and  heat  but  do  not  boil.  This  is  a  method  of 
separating  arsenic  sulfide  from  many  other  sulfides.  To  the  solution 
add  an  excess  of  dilute  HC1  which  will  reprecipitate  the  arsenic  sulfide  • 
together  with  some  sulfur. 

83. — Treat  dilute  solutions  of  sodium  arsenite  and  arsenate  with 
silver  nitrate  and  note  carefully  the  colors  of  the  precipitates.  Try 
the  solubility  of  portions  of  each  with  dilute  nitric  acid  and  with  am- 
monia. Treat  a  solution  of  the  sodium  arsenate  with  ammonium  chlo- 
ride, make  alkaline  with  ammonia  and  add  a  solution  of 'a  magnesium 
salt.  Compare  the  precipitate  with  that  obtained  in  (82)  in  the  same 
way  and  also  under  phosphoric  acid.  How  may  an  arsenate  be  dis- 
tinguished from  a  phosphate?  Given  a  mixture  of  arsenate  and  phos- 
phate how  could  you  remove  the  AsO4  and  test  for  the  PO4? 

84. — Tests  for  Small  Amounts  of  Arsenic:  One  of  the  following 
may  be  used: 


44 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


Fig.  16. 


(a)  The  Modified  Gutzeit  Test:  Good  filter 
paper  should  be  wet  with  a  solution  of  mercuric 
chloride,  allowed  to  dry,  and  cut  into  strips 
narrow  enough  to  slip  into  the  horizontal  tube,  fig. 
16.  An  arsenic  solution  containing  0.01  mg.  per 
c.c.  and  another  containing  1  mg.  per  c.c.  is  made 
up  for  class  use.  On  no  account  should  ordinary 
arsenic  solutions  be  used  in  this  or  the  next,  (b). 

Place  a  strip  of  the  pap*kr  in  the  tube,  and  about 
10  grams  of  coarsely  granulated  zinc  in  the  bottle, 
and  add  15  c.c.  of  dilute  HC1  with  an  equal  volume 
of  water.  After  a  little  time  look  to  see  whether 
the  paper  shows  by  brownish  color  that  the  mater- 
ials used  contain  arsenic,  and  if  not  add  5  c.c.  of 
the  arsenic  solution  of  0.01  mg.  of  arsenic.  Let 
stand  20  minutes  and  observe  color  of  paper. 

(b)  Set  up  apparatus  as  in  fig.  17,  the  outer  section  of  delivery 
tube  being  of  hard  glass.  Place  in  flask  10  grams  pure  zinc  and  add 

pure  dilute  sulfuric  acid. 
When  air  is  expelled  light 
the  gas  with  a  test  tube. 
Now  heat  the  tube  and  add 

f~7  5  c.c.  of  the  solution  mark- 

^  ed  .01  mg  of  arsenic  to  1 

c.c.  After  several  minutes 
examine  arsenic  mirror 
beyond  the  heated  por- 
tion. Move  flame  toward 
mirror.  Is  the  latter  easily 
volatilized?  Now  remove 
lamp  and  add  to  flask  5 
c.c.  of  the  solution  marked 
1  mg.  of  arsenic  to  1 
c.c.  Hold  in  the  flame  a 

porcelain  evaporating  dish 

Fig.  17.  or  crucible  lid.      Examine 

the  spots  of  arsenic  form- 
ed, and  try  the  action  of  a  solution  of  sodium  hypochlorite  upon  them. 
Antimony  also  gives  a  similar  mirror  and  spots,  but  it  is  not  soluble  in 
the  hypochlorite. 

85. — If  desired  to  use  either  of  the  tests  for  arsenic  for  practical 
purposes,  prepare  five  standard  papers  or  five  standard  tubes  con- 
taining .01,  .02,  .03,  .04  .05,  mg.  of  arsenic  by  using  1,  2,  3,  4  and  5  c.c. 
of  the  solution  containing  .01  mg  to  1  c.c. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  45 

Now  cut  a  piece  of  highly-colored  wall-paper  for  example,  one 
decimeter  square,  tear  into  bits,  place  in  a  dish,  heat  with  5  c.c.  con. 
sulfuric  acid,  until  the  paper  is  completely  charred.  Let  cool,  treat 
contents  of  dish  with  25  c.c.  of  water,  stir  well  and  filter.  Test  the 
solution  for  As  as  above,  using  only  so  much  as  necessary  to  obtain  a 
mirror  or  paper  containing  .01  to  .05  mg.  of  arsenic.  Compare  with 
standards,  and  calculate  the  amounts  of  As  to  one  sq.  meter  of  wall 
paper.  A  similar  proceeding  will  suffice  for  cotton  fabric. 

AtfTIMOXY. 

86. — Heat  about  0.5  gram  of  finely  powdered  antimony  with  about 
10  c.c.  con.  HC1.  Does  it  dissolve?  Now  add  to  the  test  tube  about  1 
c.c.  of  con.  HNO3,  being  careful  that  the  liquid  does  not  froth  over  on 
the  hand.  When  most  of  the  metal  is  dissolved  pour  a  little  of  the  so- 
lution into  test  tube  of  water,  and  another  portion  into  a  solution  of 
salt.  Compare  the  precipitate  of  antimony  oxy-chloride  SbOCl  in  the 
two  tubes,  allowing  for  salt  which  may  precipitate  in  the  second  tube 
and  quickly  settle.  Add  con.  HC1  to  the  first  tube  till  the  precipitate 
just  dissolves,  and  pass  an  excess  of  H2S  into  the  two  solutions.  Also 
pass  the  gas  into  solutions  of  tin  chloride  and  bismuth  chloride.  Fil- 
ter off  the  precipitated  antimony  sulfide  and  wash  on  filter.  Remove  a 
little  of  the  precipitate  Sb2S3  to  a  dish  and  boil  with  a  little  con.  HC1. 
How  may  antimony  be  separated  from  arsenic?  See  82. 

Dissolve  a  portion  of  the  sulfides  of  antimony  and  tin  by  warming 
in  a  dish  with  yellow  ammonium  sulfide,  and  try  to  dissolve  bismuth 
sulfide  in  the  same  way.  To  the  solutions  of  Sb  and  Sn  sulfides  add  an 
excess  of  dilute  HC1  when  the  sulfides  will  be  reprecipitated.  How 
may  antimony,  arsenic  and  tin  in  solution  be  separated  from  bismuth 
and  most  other  heavy  metals? 

To  illustrate  the  separation  of  antimony  and  tin  dissolve  their  sul- 
fides together  in  hot  con.  HC1,  boil  for  a  few  moments,  dilute  the  so- 
lution with  twice  its  volume  of  water,  put  in  "card  teeth"  or  steel  wool 
and  boil  persistently.  Filter  off  the  black  scales  of  Sb  and  to  the  fil- 
trate'add  a  solution  of  mercuric  chloride.  A  white  precipitate  of  mer- 
curous  chloride  shows  stannous  chloride,  SnCl2,  was  present.  Dissolve 
the  black  scales  on  the  filter  with  a  very  little  aqua  regia,  dilute  with 
20  times  its  volume  of  water  and  pass  in  H2S  which  will  precipitate 
antimony  sulfide. 

CARBOff. 

87. — In  a  test-tube  held  in  a  clamp  with  its  mouth  slightly  lower 
than  the  other  end,  heat  a  block  of  wood.  Light  the  gas  given  off  and 
let  any  liquid  distilled  fall  into  a  vessel  not  worth  cleaning.  Examine 


46  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

liquid  and  test  for  acid.  When  the  volatile  products  cease  to  come  off 
examine  the  charcoal  left  in  tube  as  to  its  form,  structure,  lightness. 

In  a  small  flask  boil  solutions  of  litmus,  cochineal,  and  indigo  with 
powdered  bone  charcoal  and  filter.  Use  about  3  g.  fresh  charcoal  with 
25  c.c.  of  each  solution.  If  any  one  of  the  filtrates  is  not  clear  return 
it  to  the  flask,  boil  again  and  filter. 

90. Preparation  of  Carbon  Dioxide:  Provide  a  bottle  or  flask 

with  thistle  tube  and  two-piece  delivery  tube  as  in  68.  Put  in  25-50 
grams  of,  calcium  carbonate,  cover  it  and  the  tip  of  thistle  tube  with 
water,  and  add  a  little  con.  HC1  from  time  to  time  to  maintain  a  mod- 
erate flow  of  COv.  Three  or  four  jars  of  gas  may  be  collected  in  the 
same  way  as  chlorine. 

Now  pass  the  gas  into  a  test  tube  half  full  of  clear  lime  water  till 
the  precipitate  of  normal  calcium  carbonate,  CaCOs,  formed  at  first 
completely  dissolves  as  it  is  changed  into  the  acid  carbonate,  H2Ca 
(CO3)2.  Reverse  the  reaction  by  boiling  one  half  of  the  solution.  To 
the  other  lialf  add  lime  water  which  will  'precipitate  the  normal  car- 
bonate. (See  116). 

Pass  CO2  into  5  c.c.  3-normal  sodium  hydroxide,  slowly  enough  to 
make  absorption  apparent.  Na2CO3  is  formed  at  first,  but  later  the  acid 
carbonate,  HNaCOs,  which  being  sparingly  soluble,  is  precipitated.  Re- 
verse the  reaction  to  the  Na2CO3  by  boiling.  Test  a  few  drops  of  the 
solution  and  also  small  amounts  of  other  carbonates  for  carbon  diox- 
ide thus :  Place  a  little  of  the  carbonate  in  one  test  tube  'and  a  few  c.c. 
clear  lime  water  in  a  larger  one.  Add  an  excess  of  dil.  HC1  to  the  car- 
bonate a,nd  pour  the  gas  set  free,  but  no  liquid,  into  the  lime  water, 
shielding  the  mouths  of  the  tubes  from  air  currents  as  shown  in  lec- 
ture room.  A  white  precipitate  in  the  lime  water  shows  carbon  diox- 
ide and  a  carbonate. 

91. — Introduce  into  a  jar  of  carbon  dioxide  rapidly  burning  phos- 
phorus. Place  a  bit  of  burning  candle  in  a  beaker  and  pour  carbon 
dioxide  upon  it  from  a  jar.  Which  is  more  dense,  CO2  or  air? 

Determine  whether  a  candle  v/ill  burn  in  the  exhaled  breath  by 
filling  the  lungs  to  capacity,  exhaling  through  a  tube  into  a  jar,  and 
without  delay  lowering  into  it  a  bit  of.  burning  candle. 

92. — Carbon  dioxide  may  be  prepared  by  heating  acid  sodium  car- 
bonate, "soda,"  or  magnesite,  MgCO*  as  in  Fig.  13  but  turning  the  de- 
livery tube  downward.  Soda  decomposes  into  the  normal  carbonate 
so  easily  that  a  test  tube  may  be  used.  For  magnesite  the  iron  tube 
is  better.  In  this  case  the  reaction  is  similar  to  that  in  making  lime 
by  heating  CaCO3. 

93. — (a)  Carbon  Monoxide:  Set  up  the  apparatus  as  in  Fig  18, 
having  the  bottles  two-thirds  full  of  NaOH,  one  part  of  the  solid  to 
three  of  water.  Place  in  flask  20  grams  oxalic  acid  and  40  c.c.  con. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 


47 


HiS04.    Prepare  a  tube  with  copper  oxide  as  in  Fig  5,  if  (b)  is  assigned. 

Heat  the  flask  so  as  to  maintain  a  stream  of  gas  slow  enough  to 
permit  the  absorption  in  the  bottles  of  the  CO2  which  comes  off  with 
the  CO. 


Fig.  18. 

After  a  jar  full  of  gas  has  come  over,  place  delivery  tube  in  a  test-tube 
of  lime  water.  It  should  not  become  turbid.  If  turbid,  decrease  the 
rate  of  flow.  Collect  two  jars  full  of  gas.  Remove  cover  of  one  jar, 
quickly  pour  in  lime  water,  seal  and  shake.  Burn  the  gas  in  the  jar 
and  shake  again.  Apply  a  flame  to  the  mouth  of  the  other  jar,  add 
lime  water,  seal  and  shake.  What  is  formed  by  burning  CO? 

Disconnect  flask  from  bottles  and  place  its  delivery  tube  in  a  test- 
tube  of  lime  water.  What  other  gas  is  mixed  with  the  carbon  monox- 
ide? To  determine  the  amount  of  this  gas  fill  a  slender  test-tube  with 
the  mixture  by  displacement  of  air,  slowly  withdraw  the  delivery  tube 
and  cover  mouth  of  test  tube  with  the  wet  finger.  Place  mouth  of  tube 
under  sodium  hydroxide  in  porcelain  dish  and  let  it  remain  until  the 
volume  of  gas  no  longer  decreases.  Estimate  the  volume  of  remaining 
gas  as  compared  with  the  original  volume  of  the  mixture  and  test  it 
with  a  flame. 

(b)  Disconnect  the  delivery  tube  and  substitute  for  it  the  tube  pre- 
pared with  copper  oxide.  Support  this  tube  on  a  ring,  connect  with  a 
test  tube  of  lime  water.  Heat  the  copper  oxide  and  pass  over  it  a  slow 
stream  of  carbon  monoxide.  Use  more  oxalic  acid  in  the  generating 
flask  if  necessary.  Is  the  copper  oxide  reduced?  What  is  formed  by 
its  action  on  carbon  monoxide?  For  the  great  importance  of  this  pro- 


48  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

perty  of  carbon  monoxide  see  the  chemistry  of  the  blast  furnace  in  the 
text  book. 

94. — Methane,  Marsh  Gas:  Mix  intimately  in  a  mortar  10  grams  of 
fused  sodium  acetate  CH3C02Na  and  10  grams  of  soda  lime.  Place  in  a 
retort  of  gas  pipe,  provided  with  a  two-piece  delivery  tube  and  heat  as 
in  the  preparation  of  oxygen,  collecting  two  jars  of  gas.  Light  the  gas 
and  study  the  degree  of  luminosity  of  the  flame.  Hold  the. flame  under 
the  mouth  of  a  dry  bottle  or  jar  and  note  appearance  of  moisture.  Put 
lime  water  in  the  jar  and  shake.  Pass  the  gas  through  lime  water  in 
a  test  tube.  What  evidence  have  you  that  methane  contains  hy- 
drogen and  carbon? 

To  the  other  jar  of  CH4  add  a  little  bromine  water  and  shake.  Does 
methane  absorb  bromine? 

95. — Acetylene:  Cuprous  chloride  is  required,  and  should  be 
freshly  prepared  as  in  126.  For  the  prepaartion  of  acetylene  fit  up  the 
apparatus  in  Fig.  10,  but  substitute  a  small  flask  for  the  test  tube. 
Place  in  flask  about  10  grams  of  calcium  carbide,  and  by  means  of  the 
funnel  and  clamp  let  in  upon  the  carbide  water,  a  few  drops  at  a  time. 
When  the  air  has  been  expelled  as  shown  by  collecting  and  burning  a 
test  tube  of  the  acetylene,  collect  three  bottles  of  the  gas.  Insert  the 
delivery  tube  into  a  solution  of  cuprous  chloride  made  alkaline  with 
ammonium  hydroxide.  The  red  precipitate  is  cuprous  acetylide,  Cu2C:>. 
In  the  same  way  precipitate  silver  acetylide  from  a  solution  of  silver 
nitrate  made  alkaline  with  ammonia.  These  acetylides  are  explosive 
and  should  be  washed  into  the  sink  before  they  become  dry. 

Insert  the  delivery  tube  into  a  test  tube  half  full  of  bromine  water 
and  let  the  gas  run  for  several  minutes.  Also,  shake  in  one  of  the  bot- 
tles a  little  bromine  water  with  acetylene.  Compare  its  action  on  bro- 
mine to  that  of  methane. 

Uncover  one  bottle  of  gas,  wait  a  moment,  then  apply  a  flame.  Why 
is  there  a  bright  flash  followed  by  quiet  combustion  giving  a  very 
smoky  flame? 

96. — Flame.  Regulate  a  Bunsen  burner  so  as  to  produce  a  small 
slightly  luminous  flame.  Observe  the  three  parts,  lower  and  inner  por- 
tion which  is  non-luminous,  luminous  portion,  outer  non-luminous  por- 
tion. Compare  with  a  candle  flame.  With  a  short  piece  of  glass  tub- 
ing draw  off  gas  from  lower  part  of  flame  and  burn  it  at  the  end  of 
the  tube.  Try  the  same  with  candle  flame.  Hold  a  piece  of  wire  gauze 
in  candle  flame  just  above  the  wick  and  observe  interior.  Can  you  light 
unburned  gas  above  the  gauze?  Hold  a  piece  of  paper  for  an  instant 
in  the  same  position  and  examine  under  surface.  Repeat  experiments 
with  gauze  and  paper  using  the  flame  of  a  Bunsen  burner.  Also  turn 
on  gas,  hold  gauze  close  to  the  burner  tube  and  light  gas  above  gauze. 
Now  move  gauze  to  one  side. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  49 

97. — Oxidizing  and  Reducing  Flames.  Regulate  a  burner  so  as  to 
produce  a  small  luminous  flame,  and  find  by  trial  the  best  position 
for  tip  of  blowpipe  to  produce  a  long  slender  blowpipe  flame  with 
well-marked  inner  blue  reducing  and  the  outer  non-luminous,  oxidizing 
portions. 

Make  a  cavity  with  butt  of  pliers  or  a  small  coin  in  a  piece  of  char- 
coal, place  in  it  a  small  quantity  of  lead  oxide  and  heat  in  the  reduc- 
ing flame  of  the  blowpipe.  Continue  until  a  globule  of  lead  remains. 
Caution:  After  using  charcoal  with  the  blowpipe  always  extinguish 
any  fire  by  holding  it  under  the  faucet  and  then  return  it  to  the  char- 
coal tray. 

98. — Fermentation  of  Glucose.  In  a  500  cc.  flask  dissolve  25  grams 
of  glucose,  1  gram  each  of  sodium-potassium  tartrate,  ammonium  nitrate 
and  soclic  phosfate  in  about  250  cc.  of  hydrant  water.  Rub  with  a  little 
water  in  a  mortar  one-tenth  of  a  cake  of  yeast  and  wash  into  flask. 
Connect  the  flask  with  a  gas  washing  bottle  nearly  full  of  clear  lime 
water,  set  in  a  warm  place  and  let  remain  three  or  four  days.  Observe 
from  time  to  time  the  carbon  dioxide  given  off.  Take  out  a  drop  of  the 
liquid,  place  on  a  slide,  cover  and  examine  for  cells  of  the  yeast  plant 
with  a  microscope. 

Arrange  the  flask  as  in  fig.  6  and  distill  off  20  c.c.  keeping  dis- 
tillate well  cooled.  Clean  the  flask,  put  in  the  20  c.c.  and  distill  off 
about  5  c.c.  Test  this  second  distillate  for  alcohol,  first  testing  for  it 
in  a  known  solution  as  follows:  to  about  5  c.c.  water  add  about  1  c.c. 
alcohol  and  a  few  crystals  of  iodine,  and  shake.  Warm  the  tube  and 
add  gradually  a  solution  of  sodium  carbonate  till  the  iodine  disappears, 
lodoform  will  appear,  at  any  rate  on  cooling.  Let  a  few  drops  cool  on 
a  watch  glass  and  examine  with  the  low  power  of  a  microscope.  .  Now 
examine  for  alcohol  the  distillate  above. 

99. — Reducing  Power  of  Sugars.  Invert  Sugar.  Fehling's  solution 
is  a  mixture  of  equal  volumes  of  two  solutions,  one  containing  34.639 
grams  of  copper  sulfate  to  500  c.c.  and  the  other  173  grams  sodium  pot- 
assium tartrate  and  60  grams  sodium  hydroxide  in  500  c.c.  The  cop- 
per in  1  c.c.  of  the  mixture  is  reduced  by  0.005  gram  of  glucose. 

To  about  10  cc.  Fehling's  solution  add  enough  glucose  to  precipi- 
tate all  the  copper  on  boiling  as  cuprous  oxide,  Cu2O,  which  will  settle, 
leaving  the  clear  liquid  above.  In  a  fresh  mixture  try  a  little  milk 
sugar.  Try  cane  sugar. 

Now  dissolve  about  a  gram  of  cane  sugar  in  100  c.c.  of  water,  add  a 
few  drops  of  con.  HC1,  heat  to  70  degrees  and  let  cool.  Boil  10  c.c.  of 
this  solution  of  invert  sugar  with  20  c.c.  of  Fehling's  solution. 

100. — Esters,  Acetic  Ester:  Acids  act  upon  alcohols  forming  esters 
and  water  much  as  they  act  upon  bases  forming  salts 


50  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

and  water.  Acetic  acid  and  alcohol  act  thus:  CH3CO2H+C2H5OH= 
CH3CO2C2HS+H2O. 

Use  the  apparatus  shown  in  Fig.  6.  With  the  delivery  tube  in  the 
test  tube  pour  into  test  tube  20  c.c.  water  and  mark  the  level  with 
gummed  paper.  Remove  the  water.  Place  in  the  flask  15  c.c.  glacial 
acetic  acid,  15  c.c.  alcohol  and  5  c.c.  con.  sulfuric  acid.  Warm  the  flask 
with  a  very  small  flame  so  that  only  a  few  drops  shall  distill  over  in 
10  minutes.  Now  increase  the  heat  and  distill  to  the  20  clc.  mark.  In- 
stead of  measuring  the  distillate  a  two  hole  stopper  and  thermometer 
may  be  used,  and  the  distillation  continued  till  it  reads  95°.  Its  bulb 
should  be  in  the  vapor  above  the  liquid. 

The  ester  contains  acetic  acid  and  alcohol.  To  remove  most  of 
these  shake  it  persistently  with  twice  its  volume  of  water  in  a  small 
flask  covered  with  the  thumb.  On  standing  the  ester  rises  and  forms 
the  top  layer.  Separate  it  from  the  water  by  using  a  separatory  fun- 
nel, or  the  funnel  with  pinch  cock  and  exit  tube  as  shown  in  Fig.  10. 
Measure  the  purified  ester. 

In  a  test  tube  shake  about  2  c.c.  of  your  ester  with  an  excess  of 
NaOH  till  it  all  disappears;  that  is,  till  it  is  all  "saponified"  forming 
sodium  acetate  and  alcohol. 

Note  odor  of  acetic  ester.  Upon  this  is  based  a  good  test  for  ace- 
tic acid  or  its  salts.  To  a  little  solid  sodium  acetate  in  a  test  tube  add 
a  few  drops  of  alcohol  and  a  few  of  con.  sulfuric  acid  and  note  odor. 
Another  test  is  this:  To  a  neutral  solution  of  an  acetate  add  3  drops 
of  ferric  chloride.  A  red  color  should  be  obtained,  and  on  boiling  a 
brown  flocculent  precipitate. 

101. — Saponification,  Soap  Making:  In  a  dish  or  beaker  place 
five  grams  of  olive  oil  or  lard  or  tallow,  add  15  c.c.  of  alcohol  and  5  c.c. 
of  sodium  hydroxide  solution  of  concentration  one  to  four.  Do  not  use 
a  pipette  for  NaOH.  With  occasional  stirring  heat  on  a  water  or  steam 
bath,  for  at  least  half  an  hour,  but  better,  an  hour.  All  the  alcohol 
and  most  of  the  water  should  be  evaporated.  Examine  the  soap  when 
cold.  If  the  saponification  is  complete  a  small  portion  should  entire- 
ly dissolve  in  water,  save  a  slight  milkiness.  Dissolve  more  of  the 
soap  and  determine  whether  it  will  form  a  lather  with  soft  water.  Try 
blowing  soap  bubbles  with  the  solution.  Determine  whether  the  feel- 
ing is  that  given  by  ordinary  soap. 

Filter  some  of  the  soap  solution.  Pour  a  little  into  distilled  water, 
into  hydrant  water,  into  solutions  of  calcium  and  magnesium  sulfates. 
Shake  each  tube  and  determine  in  which  ones  a  persistent  lather  is 
formed.  To  those  in  which  it  is  not  formed  add  more  soap  solution  and 
shake  again,  until  a  lather  persists.  Calcium  and  magnesium  com- 
pounds make  water  "hard."  Why  do  such  waters  require  more  soap? 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  51 

To  a  concentrated  solution  of  soap  add  dilute  hydrochloric  acid  in 
excess.  The  precipitate  consists  of  a  mixture  of  organic  acids  of 
which  stearic  acid,  Cn  H33COOH,  is  representative.  The  oils  and  fars 
mentioned  above  consist  of  mixtures  of  what  are  known  as  esters  of 
these  acids.  The  most  common  perhaps  is  stearin,  (Ci-HusCOOhCsHs. 
Write  the  equation  for  the  saponification  of  stearin  with  sodium  hy- 
droxide; also,  write  equation  for  action  of  soap  on  solution  of  calcium 
sulfate. 

SILICON. 

102.— Silicic  Acid,  To  10  c.c.  of  a  solution  of  sodium  silicate  in  a 
small  beaker  add  con.  HC1  drop  by  drop  stirring  for  a  few  moments 
after  each  addition.  Silicic  acid  will  separate  as  a  jelly.  Collect  the 
silicic  acid  on  a  filter,  wash  with  water,  transfer  to  a  crucible,  dry 
over  flame  and  finally  ignite.  When  cold  try  to  dissolve  the  residue 
in  water,  in  hydrochloric  acid. 

103.— Boric  Acid:  Place  15  g.  borax  in  50  c.c.  water  in  a  beaker. 
Heat  till  dissolved  and  add  15  c.c.  con.  HC1  with  stirring  and  allow  the 
liquid  to  cool  thoroughly.  Filter  off  and  examine  the  boric  acid.  Dis- 
solve a  little  of  the  substance  in  alcohol  in  a  dish,  set  fire  to  alcohol 
and  observe  color  of  the  flame.  Moisten  a  strip  of  turmeric  paper  with 
the  solution  from  the  boric  acid.  Observe  color,  then  treat  the  paper 
with  a  solution  of  sodium  carbonate,  let  dry,  best  on  steam  bath,  and 
note  color. 

QUALITATIVE  TESTS  FOR  THE  COMMON  ACIDS 

106. — The  substance  given  for  analysis  may  be  in  solution.  If  not, 
dissolve  in  water,  heating  if  needed.  If  not  soluble  in  hot  water  use  di- 
lute nitric  acid,  and  heat.  If  gas  is  given  off  one  or  more  acids  of 
group  (1)  are  present,  and  the  special  tests  may  be  applied  at  once. 

Group  (1) :  To  a  small  portion  of  the  solution  add  dilute  HNO3  in 
excess  if  not  already  added,  and  heat.  If  gas  is  set  free  try  to  identify 
it  by  odor,  color  anft  tests.  Only  CO2,  S02,  H2S,  NO-,  are  likely  to  occur. 
Make  special  tests  for  the  following  acids  as  given  in  the  sections  in- 
dicated by  the  numbers:  H2CO3  (90),  H2S03  (71),  H2S2O3  (73),  H2S  (68), 
HN02  (54b). 

Group  (2) :  To  another  portion  of  the  solution  add  dilute  nitric 
acid  in  excess  if  not  already  present  and  if  members  of  group  (1)  are 
present  boil  to  expel  any  gas.  To  a  portion  of  the  boiled  solution  add 
barium  chloride  in  excess.  A  white  precipitate  shows  H2SO4  (71). 
Filter  it  off  and  make  nitrate  alkaline  with  ammonia.  A  yellow  pre- 
cipitate indicates  H2CrO4  (140).  A  white  precipitate  indicates  one  or 
more  of  the  following:  H3PO.,-(80),  H3AsO4  .  (83),  HsAs03  (83), 
H3BO3  (103),  HF  (38),  H2C2O4  (see  following):  To  another  portion  of 


52  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

the  boiled  solution  add  ammonia  in  excess,  then  CaCl2.  An  immediate 
precipitate  may  mean  any  of  the  acids  of  this  group  save  sulfuric.  Add 
an  excess  of  acetic  acid.  If  the  precipitate  is  insoluble,  oxalic  acid  or 
hydrofluoric  or  both  are  present.  To  a  third  portion  of  the  boiled  solu- 
tion add  manganese  dioxide,  which  will  give  C02  if  oxalic  acid  is  pres- 
ent. It  is  well  in  any  case  to  make  the  special  test  for  boric  acid,  since 
its  barium  and  calcium  salts  are  distinctly  soluble. 

Group  (3)  To  the  original  solution  add  silver  nitrate.  If  a  pre- 
cipitate is  formed  add  an  excess  of  dilute  nitric  acid.  If  the  precipi- 
tate all  dissolves  this  indicates  some  acid  in  previous  groups.  If  in- 
soluble, one  or  more  of  the  following  are  present:  HC1  (37),  HBr  (37), 
HI  (37),  H4Fe(CN)6  (142),  HsFe(CN)6  (142).  A  pure  white  precipitate 
shows  only  HC1.  If  the  precipitate  is  colored  proceed  to  the  special 
tests  for  the  others.  For  halogen  acids  see  also  127b. 

Group  (4)  Nitric  and  acetic  acids  must  always  be  tested  for  if 
the  substance  is  soluble  in  water.  For  HNO3  see  (56b) ,  and  for  acetic, 
(100). 

SODIUM,  POTASSIUM,  LITHIUM. 

107. — Through  a  spectroscope  suitably  adjusted  by  the  instructor 
examine  a  flame  colored  by  a  sodium  salt,  and  locate  the  sodium  line  on 
the  scale.  Locate  in  the  same  way  the  red  line  given  by  potassium, 
and  that  given  by  lithium.  Draw  a  millimeter  scale  in  your  note  book 
and  place  the  lines  in  their  correct  positions.  If  the  spectroscope  has 
no  scale,  estimate  as  closely  as  possible  the  relative  positions  of  the 
three  lines.  Note  that  any  chemical  will  show  some  sodium.  Its  line 
is  no  evidence  that  sodium  is  present  in  considerable  quantity. 

108. — What  is  the  action  of  sodium  on  water?  Standing  well  back 
drop  a  bit  of  potassium  into  water  in  a  bottle  or  beaker.  How  does  the 
result  differ  from  that  given  by  Na? 

To  reduce  the  violence  of  the  action  of  the  alkali  metals  on  water 
their  alloys  with  some  other  metal  are  often  used.  Drop  into  water 
"hydrone"  which  is  an  alloy  of  sodium  and  lead.  Also  try  sodium 
amalgam,  an  alloy  of  sodium  and  mercury.  When  the  action  of  the 
latter  is  over  put  the  mercury  in  a  dish  provided  for  the  purpose. 
Never  put  mercury  into  sinks.  Why?  Test  the  solution  in  each  case 
with  red  litmus  or  turmeric  paper.  Test  soapy  feeling  of  the  solution 
between  thumb  and  finger. 

109.— Preparation  of  Sodium  Hydroxide:  Dissolve  5  grams  of  so- 
dium carbonate  in  75  c.c.  of  water  in  a  porcelain  dish  and  reserve  5  c.c. 
Heat  the  remainder  nearly  to  boiling,  and  stir  in  a  little  at  a  time  5 
grams  of  slaked  lime.  Boil  gently  for  several  minutes,  replacing  the 
water  evaporated,  let  settle  and  filter  off  the  solution.  If  it  destroys 
the  filter,  let  cool,  add  a  little  water  and  use  another  filter.  To  prove 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  53 

the  reaction,  Na2CO3+Ca(OH)2=CaCO3+2NaOH,  proceed  as  follows: 
Wash  the  insoluble  residue  in  the  dish  twice  by  filling  nearly  full  of 
water,  stirring,  letting  the  substance  settle  and  pouring  off  the  water. 
Now  put  a  little  water  on  the  insoluble  substance  in  dish,  and  place 
about  the  same  amounts  of  lime  and  water  in  another  dish  and  pour 
upon  each  about  10  c.c.  of  dilute  HC1.  Compare  the  amounts  of  carbon 
dioxide  given  off.  Now  treat  the  5  c.c.  reserved  solution  of  sodium  car- 
bonate and  the  same  volume  of  the  filtrate  with  the  same  volumes  of 
HC1  and  compare  the  gas  given  off.  You  started  with  a  soluble  car- 
bonate and  a  very  sparingly  soluble  hydroxide,  "slaked  lime."  What 
did  you  obtain  by  the  reaction? 

110.  Comparative  Tests  of  Crude  and  Pure  Sodium  Hydroxides: 
Dissolve  1-2  grams  of  crude  and  pure  NaOH,  each  in  10  c.c.  of  distilled 
water.  Make  each  acid  throughout  with  dilute  nitric  acid  and 
warm.  If  either  gives  off  gas  test  for  carbon  dioxide  as  in  90.  Test 
small  portions  of  each  acidified  solution  for  Cl  with  silver  nitrate,  and 
other  portions  for  SO4  with  BaCl2.  Test  other  portions  for  iron  by  add- 
ing to  each  a  few  drops  of  potassium  ferrocyanide  and  potassium  ferri- 
cyanide,  which  will  give  a  blue  color  if  iron  is  present. 

HI.— Purification  of  Common  Salt:  If  practicable  use  the  crude 
rock  salt  of  the  feed  store,  but  ordinary  salt  will  do.  Dissolve  about  5 
grams  in  25  c.c.  of  water.  To  one  third  add  dilute  HC1  and  test  for 
SO4,  and  preserve  the  tube  and  contents.  To  another  portion  add  a 
little  ammonium  chloride,  make  alkaline  with  ammonium  hydroxide 
and  add  ammonium  carbonate.  The  precipitate  is  calcium  carbonate. 
Filter  it  off  and  to  the  filtrate  add  sodium  phosphate  which  will  give  a 
precipitate  on  standing  if  magnesium  is  present.  See  80.  Save  this 
tube  and  contents. 

Make  a  fully  saturated  solution  of  the  crude  salt,  first  reducing  it 
to  powder  if  rock  salt  is  used,  and  shaking  a  long  time  with  water. 
Filter  if  necessary  and  to  the  solution  add  an  equal  volume  of  pure 
con.  HC1.  This  will  precipitate  most  of  the  salt.  If  gasous  HC1  were 
added  more  salt  would  precipitate.  Filter  off  the  salt  and  wash  with 
three  small  portions  of  water.  Dissolve  some  of  this  salt  in  water, 
test  for  SO4,  calcium,  magnesium,  as  above  and  compare  with  the 
results  of  these  tests  with  crude  salt. 

112.— An  Acid  Salt,  Acid  Potassium  Tartrate,  Cream  of  Tartar: 
This  is  one  of  the  few  slightly  soluble  salts  of  potassium.  As  the  term 
is  commonly  used  there  are  no  insoluble  salts  of  Na  or  K.  Dissolve 
about  4  grams  of  pure,  dry  potassium  carbonate  in  25  c.c.  of  water  and 
10  grams  of  tartaric  acid  in  50  c.c.  of  water,  measuring  the  latter  so- 
lution. To  the  carbonate  solution  add  one  drop  of  methyl  orange,  then 
add  cautiously  the  acid  solution  till  a  faint  red  color  is  obtained.  This 
gives  the  soluble  normal  salt,  K2C4H4O6.  Note  what  volume  of  the  acid 


54  LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY 

solution  was  added,  then  add  as  much  more,  stir  and  let  stand  some 
time.  The  precipitate  .is  the  acid  tartrate,  cream  of  tartar.  Which 
salt  is  least  soluble? 

Filter  off  the  cream  of  tartar,  let  it  dry  on  the  filter  and  weigh 
with  a  balanced  filter.  Calculate  the  weight  of  acid  sodium  carbonate 
to  mix  with  it  to  make  one  kind  of  baking  powder  which  when  wet  acts 
thus:  HKC4H4O6+HNaCO3=NaKCiH4Oc+H2O+CO2.  What  makes  the 
bread  rise  when  baking  power  is  used?  Try  your  baking  powder  in  a 
tube  with  water  and  test  for  COz. 

113. — Potassium  Nitrate  from  Sodium  Nitrate:  Dissolve  in  50  c.c. 
H2O,  25  g.  sodium  nitrate  and  the  calculated  amount  of  potassium  chlo- 
ride required  in  the  reaction 

KCl+NaN03=KNO3+NaCl. 

Evaporate  to  one-half  the  volume,  let  the  separated  salt  settle,  and  de- 
cant the  clear,  hot  liquid  into  a  beaker,  press  solid  with  spatula  and 
let  liquid  run  into  beaker.  This  solution  should  turn  solid  when  cold. 
Transfer  this  to  a  filter  and  let  drain.  Press  solid  between  folds  of 
filter  paper,  dissolve  in  least  hot  water  and  let  crystallize.  Transfer 
crystals  to  filter  and  let  dry. 

114. — Qualitative  Analysis:  Determine  first  the  presence  of  the 
ammonium  radical  and  the  alkali  metals  in  known  substances,  and 
then  their  presence  or  absence  in  "unknowns,"  using  the  scheme  as 
given  152,  Group  V.  In  the  initial  work  fresh  substances  will  be  ex- 
amined, and  of  course  what  is  said  of  filtrates  from  previous  groups 
and  their  preparation  for  analysis  does  not  apply. 

BARIUM,  STRONTIUM,  CALCIUM,  (MAGNESIUM). 

115. — Upon  a  piece  of  quick  lime  drop  water  as  Icng  as  it  is  taken 
up,  place  it  in  a  dish  and  observe  from  time  to  time.  After  it  has  be- 
come powdery  place  some  of  the  "slaked"  lime  in  a  jar  of  water,  shake 
it  thoroughly  and  let  settle.  Filter  a  portion  of  the  nearly  clear  lime 
water  placing  the  funnel  in  a  flask  to  protect  from  the  carbon  dioxide 
of  the  air.  Test  the  clear  solution  with  turmeric  paper.  To  portions 
of  the  lime  water  add  one-third  of  their  volumes  of  ferric  chloride  and 
magnesium  chloride  respectively.  The  precipitates  are  hydroxides 
of  the  metals.  Test  the  alkalinity  of  barium  hydroxide  and  its  action 
on  solutions  of  the  same  metals. 

116. — Pass  carbon  dioxide  into  25  c.c.  of  clear  lime  water  until  the 
precipitate  ot  calcium  carbonate,  CaCOs,  dissolves,  forming  the  acid 
carbonate,  H.CaCCOaK  The  latter  is  the  chief  substance  that  gives 
"temporary  hard  water."  To  a  small  portion  of  the  solution  add  clear 
lime  water.  How  may  lime  soften  temporary  hard  water?  Will  lime 
also  remove  magnesium  from  water?  Boil  another  portion  of  the  so- 


LABORATORY  MANUAL,  OF  GENERAL  CHEMISTRY  55 

lution  which  will  reverse  to  the  left  the  reaction  which  occurred  with 
CO2  in  excess. 

CaC03+H20+C02--=(reversibly)H2Ca(C03)2. 

117. — Pure  Calcium  Chloride:  Pour  off  the  liquid  from  flask  in 
which  carbon  dioxide  was  made,  make  it  alkaline  with  milk  of  lime  ob- 
tained by  shaking  slaked  lime  and  water  in  a  jar  and  pouring  at  once. 
Let  settle  and  filter.  How  does  this  remove  iron  and  magnesium? 
Pass  into  the  filtrate  C02  and  boil.  Why?  Filter  again  if  necessary 
and  evaporate  to  dryness  in  a  porcelain  dish  and  heat.  Expose  a  little 
of  the  calcium  chloride  to  air  till  next  period  and  observe  again.  Is  it 
deliquescent?  Dissolve  the  remainder  in  about  10  times  it  weight  of 
water,  and  use  as  calcium  chloride  solution. 

118— Comparative  Solubilities  of  Salts  of  Ba,  Sr,  Ca,  Mg:  Carbon- 
ates: To  solutions  of  Ba,  Sr,  Ca  and  Mg  chlorides  from  shelf  add  equal 
volumes  of  water  then  to  each  about  one-fifth  of  its  volume  of  ammo- 
nium chloride,  and  finally  to  each,  ammonium  carbonate.  What  com- 
pounds are  precipitated?  How  could  magnesium  be  separated  from 
the  other  three  metals?  Add  to  its  solution  sodium  phosphate,  and 
see  the  tests  for  phosphoric  acid  (80)  arid  Mg  (128). 

119. — Chromates:  To  solutions  as  under  carbonates,  but  omitting 
Mg,  add  a  little  acetic  acid  then  a  solution  of  pure  potassium  chro- 
rhate,  or  dichromate.  How  could  barium  be  separated  from  strontium 
and  calcium? 

120.— Suit* ates :  To  solutions  of  Ba,  Sr,  and  Ca  chlorides  add  a  so- 
lution of  magnesium  sulfate  and  let  stand  for  a  few  moments.  How 
does  this  prove  that  Ba,  Sr,  and  Ca  sulfates  are  less  soluble  that  mag- 
nesium sulfate?  To  fresh  solutions  of  Ba  and  Sr  chlorides  add  a  solu- 
tion of  calcium  sulfate.  How  do  the  results  show  that  Ba  and  Sr  sul- 
fates are  less  soluble  than  calcium  sulfate?  To  a  solution  of  barium 
chloride  add  a  solution  of  strontium  sulfate  and  let  stand  a  short  time. 
How  do  we  know  that  barium  sulfate  is  less  soluble  than  strontium 
sulfate?  Arrange  the  sulfates  in  the  order  of  their  increasing  solu- 
bility in  water. 

121.— Calcium  Oxalate:  To  a  solution  of  calcium  chloride  add  an' 
excess  of  ammonium  carbonate,  let  stand  a  few  moments  and  filter.  To 
the  filtrate  and  also  to  a  solution  of  calcium  sulfate  add  a  solution  of 
ammonium  oxalate  and  let  stand  half  an  hour.  State  how  you  know 
that  calcium  oxalate  is  less  soluble  than  calcium  carbonate  or  calcium 
sulfate.  Devise  a  scheme  for  the  separation  of  Ba,  Ca,  Sr  and  Mg. 

Calcium  sulfate  is  present  in  many  natural  waters  and  causes 
"permanent  hardness"  in  the  sense  that  it  is  not  precipitated  by  boil- 
ing, though  on  concentration  by  evaporation,  it  forms  a  hard  deposit 
on  the  boiler.  To  calcium  sulfate  solution  add  a  solution  of  sodium 
carbonate.  What  two  chemicals  may  be  added  to  soften  water  showing 


56  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

both  temporary  and  permanent  hardness? 

Test  the  water  of  the  laboratory  for  both  sorts  of  hardness. 

122.— Make  analyses  of  a  solution  containing  barium,  strontium, 
calcium  and  magnesium  according  to  the  directions  of  Group  IV;  also, 
analyses  of  unknown  solutions  or  solids  which  may  contain  these  me- 
tals, and  solutions  or  solids  which  may  contain  also  metals  of  Group 
V  and  the  ammonium  radical. 

DECIHORMAL  SOLUTIONS,  VOLUMETRIC  ANALYSIS. 

123. — Weigh  accurately  a  small  dish  or  beaker,  add  2.65  grams  to 
the  weights  and  exactly  balance  with  pure  sodium  carbonate.  Dis- 
solve the  carbonate  in  water,  transfer  with  rinsings  of  dish  to  a  half 
liter  flask,  using  a  funnel  and  taking  care  that  none  of  the  solution  is 
lost.  Do  not  even  lose  some  by  removal  on  the  stirring  rod.  Make  up 
the  volume  to  the  mark  with  water,  and  mix  by  placing  the  thumb 
over  mouth  of  flask  and  inverting  several  times.  This  is  a  decinormal 
solution  of  Na2CO3.  Why?  Transfer  this  solution  to  a  bottle  or  larger 
flask  and  wash  the  graduated  flask. 

Measure  in  a  small  cylinder  7.5  c.c.  pure  con.  HC1  and  dilute  it  10 
700  c.c.  and  mix  well.  Fill  a  buret  with  the  acid  solution  to  well  above 
the  zero,  fill  the  tip  of  buret  and  bring  the  surface  to  or  below  the 
zero.  Read  accurately  at  the  lowest  point  of  the  meniscus.  With  a  pipet 
(see  4)  place  20  or  25  c.c.  according  to  capacity  of  the  pipet,  of  the  so- 
dium carbonate  solution  in  a  dish  or  beaker,  add  to  it  2  drops  of  methyl 
orange.  For  comparison  it  is  well  to  place  beside  it  about  50  c.c.  of 
water  and  add-  to  it  two  drops  of  the  indicator.  From  the  buret  run  in 
the  acid  as  rapidly  as  you  wish  to  about  15  c.c.  then  a  few  drops  at  a  time 
with  stirring  till  the  solution  becomes  faintly  red  as  shown  by  compar- 
ison with  the  indicator  in  water.  Make  another  titration,  which  should 
agree  within  a  few  tenths  with  the  first. 

Divide  the  volume  of  the  alkali  by  that  of  the  acid  which  gives  the 
decinormal  concentration  factor  of  the  acid.  Measure  500  c.c.  of  the 
acid  and  multiply  by  the  factor,  which  will  give  the  total  volume  to 
which  the  500  must  be  made  up  with  water  to  become  decinormal. 
Why?  Transfer  the  500  c.c.  to  a  larger  vessel,  add  the  necessary  water, 
mix  and  titrate  again  against  the  alkali.  They  should  neutralize  each 
other  volume  for  volume. 

With  these  two  standard  deci-normal  solutions  the  concentration 
of  any  other  acid  or  alkali  may  be  determined,  or  solutions  of  desired 
concentration  may  be  made  by  the  same  method  used  in  making  the 
acid  solution. 

124. — In  titrating  weak  acids  methyl  orange  cannot  be  used ;  see 
next  section.  One  must  use  a  very  weakly  acid  or  neutral  indica- 
tor, such  as  litmus,  phenoltalein,  congo  red.  But  these  are  affected  by 


LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY  57 

the  carbonic  acid  from  the  carbonate  and  one  must  titrate  the  solution 
at  the  boiling  point  or  use  NaOH  free  from  carbonate.  Find  the  per 
cent  of  acid  in  vinegar  by  running  it  from  a  buret  into  a  measured 
volume  of  the  carbonate  boiling  and  containing  a  few  drops  of  phenol- 
talein,  or  use  cold  NaOH  supplied  by  the  instructor,  instead  of  Na2COs. 
Methyl  orange  and  phenolphthalein  are  weak,  complicated  organic 
acids.  The  following  will  give  a  correct,  general  idea  of  their  action: 
Any  strong  acid  sets  free  any  weak  acid  from  its  salt.  Let  NaR  be 
such  a  salt  where  R  is  the  negative  radical.  Then, 


In  the  case  of  phenolphthalein  R"  gives  the  red  color,  while  HR  is 
colorless.  Phenolphthalein  is  such  a  weak  acid  that  even  carbonic 
acid  is  strong  enough  to  act  in  the  same  way  as  HC1  in  the  equation 
and  form  HR.  Hence  when  the  stage  in  the  titration  represented  by 
HNaCOs  is  passed,  and  H2COa  is  formed  this  acts  as  a  relatively  strong 
acid,  forms  HR  and  thus  destroys  the  color.  On  the  other  hand  methyl 
orange  is  a  stronger  acid  than  carbonic  and  the  yellow  color  of  its 
negative  ion  in  Na++R~  persists  till  there  is  a  slight  excess  of  HC1, 
when  red  HR  is  formed. 

125.—  The  facts  stated  in  124  are  well  illustrated  by  titrating  in  the 
same  solution  both  normal  and  bicarbonate  by  using  different  indica- 
tors, as  fallows: 

Dissolve  about  0.2  gram  of  normal  sodium  carbonate  in  25  c.c.  of 
water,  without  heating,  add  a  few  drops  of  phenoltalein,  fill  a  buret 
with  your  deci-normal  acid,  read  and  run  into  the  carbonate  solution 
drop  by  drop  near  the  end,  till  the  pink  color  just  disappears.  The 
carbonate  is  now  all  HNaCOs.  Add  a  few  drops  of  methyl  orange,  read 
the  buret  and  run  in  the  acid  till  the  solution  takes  on  a  faint  tinge  of 
red,  using  the  indicator  in  water  for  comparison  as  in  123.  Read  again 
and  compare  the  volumes  required  to  change  the  normal  to  the  bicar- 
bonate, and  to  neutralize  the  latter.  '  Why  are  they  approximately 
equal  ?  Why  are  these  carbonates  alkaline,  having  no  ion  OH  ?  See  6  1. 

COPPER,  SILVER. 

Copper. 

126.  —  (a)  Dissolve  5  grams  copper  chloride  in  10  c.c.  of  con.  HC1 
and  10  c.c.  water;  or  (b)  prepare  a  solution  of  the  copper  chloride  by 
dissolving  5  grams  copper  sulfate  and  2.5  grams  common  salt  by  heat- 
ing with  10  c.c.  of  water  in  a  test  tube.  When  dissolved  set  the  tube 
in  cold  water  for  several  minutes.  Pour  off  the  solution  from  the  sep- 
arated sodium  sulfate.  Why  is  this  formed?  Now  add  to  the  solution 
10  cc.  con.  HC1,  let  stand  a  few  moments  and  filter  off  the  salt.  Why 
is  salt  thus  formed?  Whether  (a)  or  (b)  boil  very  gently  the  solu- 


58  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

tion  with  5  grams  finely  divided  copper  in  a  small  flask,  replacing 
evaporated  liquid  if  necessary  with  dil.  HC1.  Continue  till  colorless 
or  till  a  few  drops  poured  into  water  gives  no  blue  color.  You  now 
have  HCuCl2.  Pour  a  part  into  water  which  decomposes  it  giving  in- 
soluble CuCl.  Does  this  dissolve  in  ammonia?  Compare  with  silver 
chloride.  Put  the  ammonia  solution  in  a  white  dish,  stir  and  note  that 
the  cuprous  ion  is  rapidly  oxidized  to  cupric  ion  as  shown  by  increas- 
ing blue  color. 

To  a  portion  of  the  liquid  from  flask  add  an  excess  of  NaOH.  What 
is  the  red  compound?  See  Fehling's  solution  (99).  Heat  a  little  of  the 
white  precipitate  with  Br  water  and  give  result  in  terms  of  ion  formed. 
Expose  some  of  the  white  CuCl  on  the  filter  to  sunlight  and  note  result 
after  an  hour. 

There  is  no  cupric  but  only  cuprous  iodide.  To  a  few  c.c.  copper 
sulfate  solution  add  a  little  potassium  iodide  solution,  and  test  the  so- 
lution for  free  iodine  with  starch  paper.  Add  Ic.c.  carbon  disulfide 
shake  and  let  settle.  Note  color  of  the  carbon  disulfide. 

Determine  whether  dilute  sulfuric,  hydrochloric  and  nitric  acids 
give  H  by  their  action  on  copper  and  cadmium  and  explain  results. 

(c) — Tests  for  Copper:  To  half  a  test  tube  of  water  add  a  drop  of 
solution  of  any  cupric  salt,  and  make  alkaline  with  ammonia.  The 
blue  color  is  due  to  the  complex  ion  Cu(NH3)4++.  To  a  like  dilute  solu- 
tion of  copper  add  a  few  drops  of  acid  and  a  little  potassium  ferrocy- 
anide.  Dilute  equal  parts  of  each  solution  till  the  colors  are  just  vis- 
ible and  state  which  test  is  the  more  delicate. 

(d) — To  dilute  solutions  of  copper,  cadmium  and  zinc  preferably 
chlorides  add  H2S  in  excess.  Let  settle  and  pour  off  most  of  the  liquid 
in  each  case  and  add  equal  volumes  of  dil.  HC1.  Are  the  reactions 
with  H2S  reversible?  How  may  Cu  and  Zn  be  separated?  To  solu- 
tions of  the  same  metals  add  an  excess  of  NaOH  and  heat  to  boiling. 
For  explanation  see  127a  and  128c.  How  may  copper  be  separated 
from  the  other  two  elements? 

To  solutions  of  copper  and  cadmium  add  an  excess  of  ammonia. 
Add  to  the  copper  tube  KCN  (dangerous)  drop  by  drop  till  the  blue 
color  disappears.  Add  the  same  volume  of  KCN  to  the  tube  contain- 
ing Cd.  Now  pass  H2S  into  each  tube.  How  may  Cu  and  Cd  be  separ- 
ated? 

Silver. 

127. —  (a)  To  a  little  silver  nitrate  solution  add  drop  by  drop  am- 
monia solution  till  the  small  amount  of  silver  oxide  at  first  formed  is 
dissolved,  forming  the  complex  ion  Ag(NH3)2+.  Add  a  little  more 
ammonia  and  try  to  precipitate  silver  chloride  with  a  small  amount  of 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  59 

NaCl  solution.  Now  add  an  excess  of  dilute  nitric  acid.  What  is  pre- 
cipitated and  why? 

To  another  portion  of  silver  nitrate  solution  add  sodium  hydroxide 
solution  which  will  precipitate  mainly  Ag2O,  but  it  is  alkaline  and  be- 
haves as  though  partially  hydrated.  A  similar  copper  hydrated  oxiclo 
is  formed  by  adding  an  excess  of  NaOH  to  a  solution  of  copper  sulfate 
and  heating  to  boiling.  Try  it.  The  compound  is  Cu(OH)2.2CuO.  Add 
ammonia  to  the  tube  containing  the  silver  oxide  till  it  just  dissolves 
and  then  H2O2,  which  will  give  metallic  silver.  Prepare  the  same  sort 
of  a  solution  of  silver  oxide  in  ammonia,  that  is,  containing  the  ion 
Ag(NH3)2+,  add  1  gram  sodium  potassium  tartrate  dissolved  in  a  little 
water,  warm  the  test  tube  and  let  stand.  If  silver  is  not  deposited  on 
the  tube  warm  again.  Note  analogy  of  the  complex  silver  ammonia  ion 
to  that  of  copper,  Cu(NH3)4++  which  is  deep  blue. 

(b) — To  four  tubes  containing  silver  nitrate  add  respectively  a 
solution  of  a  chloride,  a  bromide,  an  iodide  and  to  the  fourth  drop  by 
drop  a  solution  of  KCN  (caution) ,  till  the  silver  cyanide  at  first  formed 
is  dissolved.  Treat  a  little  of  each  of  the  halides  of  silver  with  an  ex- 
cess of  ammonia.  Which  are  dissolved?  Treat  other  small  portions 
with  a  solution  of  sodium  thiosulfate  and  shake  till  dissolved.  Treat 
yet  other  portions  with  a  solution  of  KCN  till  dissolved.  How1  could 
you  distinguish  the  three  halides  of  silver  by  their  color?  How  dis- 
tinguish by  their  solubility  in  ammonia?  What  is  formed  when  sil- 
ver chloride  dissolves  in  ammonia?  What  when  it  dissolves  in  KCN? 
Expose  a  little  chloride  to  sunlight  and  observe  color  after  a  few  min- 
utes. Compare  cuprous  chloride  and  silver  chloride  as  to  the  effects 
on  them  of  sunlight  and  ammonia  solution. 

(c)  To  solutions  of  sodium  phosphate,  potassium  chromate,  sodium 
arsenite  and  sodium  arsenate  add  silver  nitrate.  Try  to  dissolve  por- 
tions of  each  precipitate  in  ammonia  and  in  dil.  nitric  acid.  Make 
careful  records  of  all  results  for  they  are  to  be  used  in  the  qualitative 
testing  for  acid  radicals. 

MAGNESIUM,  ZIJC,  CADMIUM,  MERCURY. 

128. —  (a)  To  solutions  of  magnesium,  zinc  and  cadmium  salts, 
preferably  sulfates  or  chlorides  add  an  excess  of  ammonium  hydrox- 
ide. Some  zinc  and  cadmium  hydroxide  are  precipitated  but  in  an  ex- 
cess of  ammonia  they  form  the  complex  ions  Zn(NH3)4++  and 
Cd(NH3)44+  which  are  soluble.  To  show  that  only  a  part  of  the  Mg  is 
precipitated  as  hydroxide,  filter  it  off  and  add  acid  sodium  phosphate 
to  the  filtrate  when  more  magnesium  as  NH4  MgPO4  will  be  precipi- 
tated. The  prevention  of  complete  precipitation  as  Mg(OH)2  by  am- 
monium hydroxide  is  due  to  the  necessary  accumulation  of  highly  ion- 
ized ammonium  salt  as  the  reaction  progresses.  To  show  this  first 


60  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

add  to  a  solution  of  magnesium  salt  ammonium  chloride  solution  and 
then  ammonium  hydroxide.  Also,  precipitate  Mg(OH)2  with  ammonia 
and  then  add  ammonium  chloride.  In  the  first  case  no  magnesium 
hydroxide  was  precipitated  and  in  the  second  it  dissolved. 

The  action  of  the  ammonium  salt  which  is  highly  ionized  is  to 
supply  the  common  ion  NH4  which  forces  the  already  slight  dissociation 
of  ammonium  hydroxide  to  the  left  in  NH4OH=(reversibly)NHr+OH- 
till  there  are  not  enough  OH  ions  to  form  sufficient  Mg(OH)2  to  exceed 
its  solubility  limit.  Furthermore,  the  NH4  ions  from  the  ammonium 
chloride  or  other  ammonium  salt  unite  with  the  OH  ions  associated 
with  the  Mg  in  solution  to  form  the  undissociated  NH4OH.  This  latter 
action  is  of  the  same  sort  as  in  the  neutralization  of  a  base  by  an  acid 
in  which  the  ion  H  of  the  acid  unites  with  the  OH  of  the  base  to  form 
undissociated  water. 

This  explanation  applies  to  the  solubility  of  several  other  hydrox- 
ides when  ammonium  salts  are  added  and  will  be  referred  to  later. 

Prove  that  acid  sodium  phosphate  alone  will  not  completely  pre- 
cipitate Mg  by  adding  to  a  magnesium  solution  an  excess  of  phosphate, 
filtering  and  then  adding  to  the  filtrate  ammonium  chloride  and  am- 
monia. Make  the  corresponding  salt  of  zinc,  NH4ZnPO4. 

(b)  Burn  a  little  Mg  ribbon,  place  some  of  the  oxide  on  moist 
turmeric  paper  and  state  whether  it  is  alkaline.     Burn  a  little  zinc 
dust  by  heating  and  stirring  in  an  iron  crucible,  and  determine  wheth- 
er it  has  any  alkaline  property. 

(c)  Precipitate  the  hydroxides  of  Mg,  Zn  and  Cd  by  adding  to 
their  solutions  NaOH  a  few  drops  at  a  time.    Now  add  more  NaOH  to 
half  of  each  and  determine  which  are  soluble  in  an  excess.     Dissolve 
the  other  half  of  each  hydroxide  by  adding  any  dilute  acid.     With  so- 
dium hydroxide  in  excess  zinc  and  cadmium  hydroxides  form  Na2Zn02 
and  Na2CdO2,  which  resemble  ordinary  salts  in  form  and  properties. 
The  hydroxide  of  Zn  and  Cd  and  many  other  elements  are  "ampho- 
teric";  that  is,  they  act  like  bases  toward  strong  acids,  and  like  weak 
acids  toward  strong  bases.    On  the  basis  of  the  ion  theory  explain  the 
dissolving  of  Mg(OH)2  by  HC1. 

(d)  To  solutions  of  Mg,  Zn,  and  Cd  add  H2S  for  several  minutes. 
Which  give  sulfides?     Shake  and  filter  off  half  of  each  precipitate 
and  to  the  filtrates  add  ammonia.    Did  the  H2S  completely  precipitate 
both  the  Zn  and  Cd  as  sulfides?    To  the  other  half  of  each  precipitate 
obtained  with  H2S  alone  add  dilute  HC1  and  finally  con.  HC1  if  needed 
to  dissolve  all  the  sulfides.    Are  the  reactions  of  Cd  and  Zn  salts  with 
H2S  both  reversible?    Which  one  is  most  easily  reversed?    How  may 
Zn,  Mg  and  Cd  be  separated? 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  61 

MERCURY 

129. — Place  10  grams  mercury  in  4  c.c.  con.  nitric  acid  diluted  with 
the  same  volume  of  water  and  let  stand  a  day.  Pour  off  liquid  and 
dissolve  the  solid  in  water  adding  a  little  dilute  nitric  acid,  letting  the 
metallic  mercury  remain.  This  is  a  solution  of  mercurous  nitrate,  (a) . 
Dissolve  about  2  grams  mercury  in  a  few  c.c.  con.  nitric  acid,  heat  and 
if  necessary  add  more  nitric  acid  and  heat  till  a  drop  of  the  solution  in 
water  gives  no  precipitate  with  HC1.  Now  dilute  with  about  50  c.c. 
water.  This  is  mercuric  nitrate,  (b). 

Treat  small  portions  of  (a)  and  (b)  with  NaOH  in  excess  which 
gives  Hg2O  and  HgO  and  not  the  hydroxides.  Compare  the  result  with 
that  obtained  with  silver  nitrate  and  NaOH,  with  a  cupric  compound 
and  NaOH  cold  and  after  boiling.  Try  ammonia  on  the  solutions  (a) 
and  (b).  No  oxides  or  hydroxides  are  formed  but  with  (a)  Hg  and 
Hg2N(NO3).  With  (b)  the  same  compound  but  no  free  Hg. 

Treat  a  little  of  (a)  with  dil.  HC1  and  determine  whether  the  HgCl 
is  soluble  in  an  excess  of  HC1  or  in  nitric  acid.  Try  the  action  of  am- 
monia which  gives  black  Hg.  HgNH2Cl.  How  may  mercury  in  the  mer- 
curous condition  be  separated  from  silver?  Oxidize  a  little  of  solution 
(a)  to  (b)  by  adding  Br  water  till  the  red  color  persists,  boiling  out 
excess  of  Br.  Prove  that  only  mercuric  mercury  is  present  by  adding 
dil.  HC1. 

To  small  portions  of  (a)  and  (b)  add  a  solution  of  KI  drop  by  drop 
giving  green  Hgl  and  red  HgL.  To  the  latter  add  an  excess  of  KI 
which  will  dissolve  forming  a  double  salt,  HgI2  2KI. 

To  a  very  dilute  solution  of  stannous  chloride  add  a  solution  of 
mercuric  chloride  which  gives  HgCl  while  SnCh,  stannic  chloride,  re- 
mains in  solution.  This  is  a  good  test  for  either  Hg  or  Sn.  What  must 
be  the  valence  of  each  ion  when  tested  for? 

To  (a)  and  (b),  and  to  solutions  of  Pb,  Cu,  Cd  add  an  excess  of 
H2S.  Let  settle,  pour  off  the  liquid  in  each  case  and  try  to  dissolve  the 
precipitates  by  boiling  with  dilute  nitric  acid.  How  may  mercury  be 
separated  from  the  other  metals?  Determine  whether  HgS  will  dis- 
solve in  ammonium  sulfide.  How  may  mercury  be  separated  from  As, 
and  Sb. 

TIN. 

130. — Dissolve  most  of  2  grams  of  tin  by  heating  with  10  c.c.  con. 
HC1,  best  in  a  small  flask  on  the  water  bath.  Pour  off  half  of  the  so- 
lution into  a  dish  and  add  50  c.c.  of  water  to  the  flask  and  use  the  so- 
lution as  stannous  chloride,  (a).  Heat  to  boiling  the  solution  in  the 
dish  and  add  con.  nitric  acid  a  few  drops  at  a  time  till  a  drop  in  a 


62  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

little  water  gives  no  precipitate  with  mercuric  chloride.  Add  50  c.c.  of 
water  forming  a  suitable  solution  of  stannic  chloride,  (b). 

To  small  portions  of  (a)  and  (b)  add  a  few  drops  of  NaOH,  then 
add  in  excess.  Are  these  hydroxides  "amphoteric"  ?  To  fresh  portions 
of  the  tin  solutions  add  yellow  ammonium  sulfide,  at  first  only  a  few 
drops,  then  an  excess  with  heating,  but  do  not  boil.  The  SnS  and  SnS2 
at  first  formed,  should  dissolve.  Now  add  HC1  in  excess  which  will 
reprecipitate  the  tin  sulfide,  SnS2  from  each  solution.  What  other  sul- 
fides  dissolve  in  ammonium  sulfide?  How  may  the  tin  sulfides  be  sep- 
arated from  those  of  copper,  lead,  mercury? 

To  a  little  of  (a)  add  a  solution  of  mercuric  chloride  which  will 
give  insoluble  HgCl.  Try  (b)  with  mercury  chloride.  Now  reduce  to 
the  stannous  condition  the  Sn  in  a  portion  of  (b)  by  persistent  heat- 
ing with  finely  divided  iron,  filter  and  add  mercuric  chloride.  This  is 
a  good  test  for  either  Hg  or  tin.  In  what  condition  of  oxidation  must 
each  be? 

To  a  portion  of  (a)  add  a  few  drops  of  gold  chloride  which  will 
give  colloidal  gold,  the  purple  of  Cassius,  a  good  test  for  gold. 

LEAD. 

131. — To  a  dilute  solution  of  lead  acetate  or  nitrate  add  an  excess 
of  NaCl  solution.  Before  the  lead  chloride  settles  pour  off  one  half 
and  boil,  adding  more  water  if  necessary  and  boiling  till  it  all  dis- 
solves. Let  the  other  half  stand  5-10  minutes,  filter  and  to  the  filtrate 
add  dil.  sulfuric  acid.  What  proof  here  that  lead  sulfate  is  less  soluble 
than  lead  chloride?  Filter  off  the  lead  sulfate  and  pass  into  nitrate 
H2S.  Which  is  less  soluble,  lead  sulfate  or  lead  sulfide?  From  a  dilute 
solution  precipitate  lead  chloride  and  prove  that  it  is  soluble  in  HC1  if 
added  in  large  excess.  Prove  in  the  same  way  that  lead  sulfate  is 
soluble  in  nitric  acid.  These  facts  must  be  kept  in  mind  in  the  analysis 
of  Group  II.  Try  to  dissolve  lead  chloride  in  ammonia.  Devise  a 
scheme  for  the  separation  of  Ag,  Hg  (mercurous)  and  Pb. 

132. — To  a  solution  of  lead  acetate  add  a  solution  of  sodium  carbon- 
ate till  alkaline.  The  precipitate  is  a  basic  carbonate  similar  to  "white 
lead."  Shake  a  solution  of  lead  acetate  with  PbO,  filter  and  pass 
through  carbon  dioxide  which  will  give  much  the  same  compound. 
Mix  5  c.c.  con.  nitric  acid  and  5  c.c.  water,  heat  and  add  a  little  at 
a  time  about  2  grams  red  lead,  PbaO^  The  brown  product  is  lead  di- 
oxide, PbO2.  Filter  a  few  drops  of  the  liquid  and  add  dil.  H2S04.  Is 
there  lead  in  solution?  To  show  the  oxidizing  power  of  lead  dioxide 
add  to  it  and  the  rest  of  the  liquid  a  few  drops  of  any  salt  of  manga- 
nese and  boil  persistently.  Let  the  solid  matter  settle  and  note  deep 
red  color  due  to  permanganic  acid,  HMn04. 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  63 

133.— To  20  c.c.  of  the  lead  acetate  solution  of  the  laboratory  add 
80  c.c.  pure  water,  place  in  flask  and  with  a  thread  suspend  in  it  a 
folded  strip  of  zinc,  about  5  grams,  and  let  stand  till  next  period.  Ex- 
amine the  "lead  tree,"  and  test  the  solution  for  lead  and  zinc.  Which 
metal  is  most  electropositive? 

ALUMINIUM. 

134. — Dissolve  a  few  tenths  of  a  gram  of  Al  in  con.  HC1  and  dilute 
to  a  test  tube  full,  and  use  where  A1C13  is  required  (a).  Dissolve  an- 
other small  amount  of  Al  in  a  few  c.c.  of  NaOH  warming,  which  gives 
sodium  aluminate,  Al  (ONa)3,  (b).  What  gas  is  given  off  in  each  case? 

To  a  little  of  (a)  add  ammonia  in  excess.  To  (a)  add  drop  by  drop 
NaOH  till  a  permanent  precipitate  is  obtained.  Dissolve  one-half  of  it 
by  adding  HC1  and  the  other  half  with  an  excess  of  NaOH.  Is  A1(OH)8 
amphoteric?  In  what  two  ways  may  it  ionize?  To  (b)  add  ammonium 
chloride  till  the  precipitate  is  permanent,  and  account  for  its  formation. 
How  could  you  separate  Al  and  Zn? 

To  portions  of  (a)  add  in  excess  sodium  carbonate,  and  ammonium 
sulfide  and  account  for  the  precipitation  of  Al  (OH)8  in  each  case? 

135. — By  heating  dissolve  15  grams  of  aluminium  sulfate, 
A12  (SOOs  18H2O  in  50  cc.  water,  stir  in  the  calculated  weight  of  am- 
monium sulfate  to  make  ammonium  aluminium  alum  (NH^aAMSOO* 
and  when  all  is  dissolved  filter  into  a  crystallizing  dish  while  hot.  Let 
a  drop  or  two  of  nitrate  fall  on  a  watch  glass,  let  the  water  evaporate 
and  examine  the  crystals  with  a  microscope.  Examine  the  crystals  in 
dish,  dissolve  a  few  and  test  the  solution  with  blue  litmus  paper.  Why 
are  the  solutions  of  Al  salts  acid?  See  63. 

CHKOMIUM. 

136. — Chromate  Ion  to  Chromium  Ion:  Dissolve  in  a  dish  10 
grams  potassium  dichromate  in  50  c.c.  water  with  heat  and  let  cool. 
Add  10  c.c.  con.  sulfuric  acid  and  then  a  little  at  a  time  50  per  cent 
alcohol  heating  if  necessary  to  start  the  reaction.  Continue  the  addi- 
tion of  alcohol  till  the  solution  has  a  green  color,  a  few  drops  in  much 
water  showing  no  brown.  Avoid  a  large  excess  of  alcohol.  Set  aside 
two-thirds  of  the  solution  and  examine  the  crystals  of  chrome  alum 
at  the  next  laboratory  period.  Dilute  the  other  third  with  20  times  its 
volume  of  water  and  use  it  below.  One  molecule  of  the  dichromate  is 
reduced  by  the  oxidation  of  three  molecules  of  alcohol  to  aldehyde, 
CH3COH.  With  aid  of  text  write  the  equation.  What  was  the  action 
of  H2S  and  of  SO2  on  acidified  solutions  of  K2Cr2O7? 

Treat  a  portion  of  your  solution  of  chrome  alum  with  an  excess  of 
ammonia.  How  separate  Cr  and  Zn?  Try  to  dissolve  a  little  of  the 
remaining  precipitate  with  NH4C1.  How  separate  Cr  from  Mg?  To  an- 


64  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

other  portion  of  chrome  alum  solution  add  an  excess  of  NaOH  and  boil. 
How  separate  Cr  and  Al?  Is  Cr(OH)8  amphoteric?  Try  portions  of 
your  chrome  alum  with  solutions  of  sodium  carbonate  and  ammonium 
sulfide.  In  each  case  the  hydroxide  is  precipitated  as  in  the  case  of  Al. 

137. — Chromium  Ion  to  Chromate  Ion:  To  a  few  c.c.  of  chromium 
nitrate  or  chloride  solution  in  a  dish  add  an  equal  volume  of  NaOH. 
Stir  into  it  gradually  2  grams  sodium  dioxide  and  heat  to  boiling.  Fil- 
ter if  necessary,  acidify  a  little  of  the  filtrate  with  acetic  acid  being 
sure  of  an  excess  by  testing,  and  add  barium  chloride.  To  a  little  di- 
lute solution  of  sodium  dichromate  add  acetic  acid  and  barium  chloride 
and  compare  precipitates. 

13^.— Dichromate  to  Chromate:  (a)  Dissolve  5  grams  sodium  di- 
chromate in  25  c.c.  water  in  a  dish,  add  slowly  with  stirring  NaOH  till 
the  solution  becomes  yellow.  Evaporate  till  the  salt  crystallizes  on 
cjoling. 

(fo) — Chromate  to  Dichromate:  Dissolve  10  grams  sodium  chro- 
mate  in  25  cc.  water  and  add  the  calculated  weight  of  con.  sulfuric 
acid,  that  is,  one  molecule  of  acid  to  two  of  the  chromate.  Evaporate 
till  the  dichromate  crystallizes  out  on  cooling. 

(c) — Chromium  Trioxide:  Dissolve  2  grams  potassium  dichromate 
»ii  5  c.c.  water  by  heat,  cool  till  it  begins  to  separate  then  drop  con. 
dlfuiic  acid  directly  upon  the  surface  of  the  liquid  in  test  tube  till  the 
precipitate  formed  does  not  quite  dissolve.  Heat  to  dissolve  most  of 
it,  note  the  escape  of  some  oxygen,  set  in  rack  to  cool  and  observe 
crystals  of  CrO3  in  an  hour. 

139.— Chromium  Oxychloride:  In  a  retort  place  3  grams  dichro- 
mate, 2  grams  NaCl  and  15  cc.  con.  H2SO4.  Heat  and  collect  the  oxy- 
chloride  as  you  did  nitric  acid.  What  does  it  look  like?  Dissolve  some 
of  it  in  water,  add  an  excess  of  NaOH,  then  an  excess  of  acetic  acid  and 
finally  barium  chloride.  What  is  the  precipitate? 

140. — Test  for  Chromate  Ion :  To  two  small  portions  of  a  chromate 
solution  add  acetic  acid  then  barium  chloride  and  lead  acetate  respec- 
tively. To  a  neutral  solution  of  a  chromate  add -silver  nitrate.  Divide 
into  two  portions.  In  one  try  the  solubility  of  the  silver  chromate  in 
ammonia  and  in  the  other  with  dil.  nitric  acid.  A  third  test  is  the  re- 
duction of  chromate  ion  to  chromium  ion  with  change  from  yellow  or 
red  to  green  as  in  136. 

IROff,  NICKEL,  COBALT. 
Iron. 

141. — Dissolve  most  of  about  1  gram  of  card  teeth  in  con.  HC1  di- 
luted with  an  equal  volume  of  water.  Pour  one  half  in  a  dish.  Dilute 
the  other  half  to  a  test  tube  full  leaving  in  it  the  undissolved  iron.  This 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  65 

is  ferrous  chloride,  FeCl2,  solution  (a).  Heat  the  half  in  the  dish  and 
with  stirring  add  a  little  con.  nitric  acid  at  a  time  till  the  black  pre- 
cipitate at  first  formed  dissolves  and  a  drop  of  the  solution  in  a  little 
water  gives  no  blue  color  with  potassium  ferricyanide.  Dilute  to  a 
test  tube  full.  It  is  ferric  chloride,  FeCls  (b). 

Treat  small  portions  of  (a)  with  an  excess  of  ammonia,  NaOH,  so- 
dium carbonate,  H2S,  ammonium  sulfide.  The  first  and  second  give 
ferrous  hydroxide  Fe  (OH)  2,  changing  in  air  to  ferric  hydroxide  Fe  (OH)  3; 
the  third  gives  ferrous  carbonate.  Hydrogen  sulfide  has  no  effect,  but 
ammonium  sulfide  gives  black  FeS.  Treat  small  portions  of  (b)  with 
the  same  reagents  in  excess.  Ammonia,  NaOH  and  sodium  carbonate 
give  ferric  hydroxide.  H2S  reduces  ferric  to  ferrous  iron  with  the  sep- 
aration of  sulfur.  Why  does  it  not  precipitate  iron  sulfide?  Is  either 
hydroxide  amphoteric?  Try  to  dissolve  ferrous  and  ferric  hydroxides 
with  ammonium  chloride.  Compare  the  action  of  sodium  carbonate 
and  ammonium  sulfide  on  ferric  iron,  aluminium  and  chromium  solu- 
tions. Devise  ways  to  separate  iron  from  Al,  Cr,  Cu,  Mg,  As. 

142.— Treat  portions  of  (a)  and  (b)  with  solutions  of  potassium 
ferrocyanide,  potassium  ferricyanide,  ammonium  sulfocyanide  and 
tabulate  results  as  tests  for  ferrous  and  ferric  iron. 

The  following  is  a  fine  example  of  a  reversible  reaction  and  illus- 
tration of  the  influence  of  the  common  ion:  To  a  test  tube  nearly  full 
of  water  add  about  5  drops  of  ammonium  sulfocyanide  and  the  same 
amount  of  ferric  chloride.  The  reaction  is, 

FeCl3+3NH4CNS=  (reversibly)  Fe  (CNS)  3+3NH4Cl. 

Divide  the  red  solution  in  four  test  tubes.  To  one  add  more  of  the 
ferric  chloride,  to  the  second  more  of  the  sulfocyanide,  to  the  third 
ammonium  chloride,  and  compare  colors  with  that  of  the  fourth.  Re- 
fer to  62  and  to  text  book  and  explain  fully. 

143.— Double  Salts:  (a)  Dissolve  in  50  cc.  water  by'  heat  10 
grams  ferrous  sulfate  and  the  calculated  amount  of  ammonium  sulfate 
to  make  ammonium  ferrous  sulfate  and  filter  hot  into  a  crystallizing 
dish.  Examine  crystals  in  dish,  also  let  a  drop  cool  on  watch  glass 
and  use  microscope.  This  is  Mohr's  salt,  (NH4)2Fe(SO4)2  6H2O. 

(b)  To  make  iron  alum  dissolve  15  grams  ferrous  suifate  in  25  c.c. 
water  in  dish,  also  the  calculated  amounts  of  ammonium  sulfate  and 
con.  sulfuric  acid.  Now  heat  and  add  slowly  with  stirring,  con.  nitric 
acid  till  a  drop  diluted  shows  no  ferrous  iron.  Set  aside  to  crystallize, 
giving  (NH4)2Fe2(S04)424H20. 

JICKEL  AND  COBALT. 

144. — From  nickel  and  cobalt  chlorides  or  nitrates  precipitate  their 
hydroxides  with  NaOH  in  excess.  Try  to  dissolve  portions  of  the  hy- 
droxides in  excess  of  NaOH  with  heating.  Are  they  amphoteric?  Try 


66  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

to  dissolve  other  portions  of  the  hydroxides  with  ammonium  chloride. 
How  may  these  metals  be  separated  from  ferric  iron,  aluminium,  chro- 
mium? 

Try  to  precipitate  the  sulfides  of  Ni  and  Co  with  H2S,  then  with 
H2S  and  ammonia  or  with  ammonium  sulfide.  Try  to  dissolve  the  sul- 
fides with  dil.  HC1.  How  may  Ni  and  Co  be  separated  from  Zn,  Fe,  Sb 
and  the  metals  of  groups  I  and  II? 

To  1  to  2  c.c.  of  solutions  of  Ni  and  Co  add  NaOH  with  shaking  till 
the  hydroxides  are  just  permanent.  Add  to  each  tube  acetic  acid  in 
slight  excess,  then  to  each  10  cc.  of  a  solution  of  potassium  nitrite  and 
let  stand.  The  precipitate  is  K3  Co(NO2)e.  How  may  Co  be  separated 
from  Ni? 

Again  precipitate  the  hydroxides  of  the  two  metals  from  1  to  2  cc. 
of  their  solutions  avoiding  a  large  excess  of  NaOH.  Add  to  each  a  so- 
lution of  KCN  (dangerous)  till  the  precipitates  just  dissolve.  Now  add 
to  each  about  1  c.c.  of  NaOH  and  bromine  water  till  it  colors  them  per- 
manently red.  A  black  precipitate  of  nickelic  hydroxide  should  be  ob- 
tained. 

MANGANESE 

145. — To  a  few  cc.  of  a  solution  of  manganous  salt,  as  MnCl2,  add 
ammonia  in  excess  and  to  another  portion  NaOH  in  excess.  Does  the 
Mn(OH)2  redissolve?  How  may  Mn  be  separated  from  Zn,  Al,  Cu,  Ag? 

Add  to  one  fresh  portion  of  the  solution  ammonia  in  excess  then 
ammonium  chloride,  and  to  another  portion  add  ammonium  chloride 
then  ammonia.  Compare  the  results  with  those  obtained  with  Mg  and 
the  same  reagents.  To  one  of  these  solutions  add  ammonium  sulfide 
and  to  the  other  hydrogen  sulfide.  How  could  you  separate  Mn  and 
Mg?  Try  to  dissolve  the  sulfide  in  dil.  HC1.  How  may  Mn  be  separ- 
ated from  Cu,  Hg,  As,  Sb? 

To  a  manganese  solution  add  ammonium  chloride,  ammonia,  and 
sodium  phosphate.  Compare  result  with  the  action  of  these  reagents 
on  solutions  of  Zn  and  Mg. 

146.— Permanganic  Acid  and  Permanganate:  To  about  2  c.c.  of 
a  solution  of  any  manganous  salt  add  an  equal  volume  of  con.  nitric 
acid,  then  about  1  gram  of  red  lead  or  lead  dioxide,  and  heat  some  time 
at  the  boiling  point.  Let  the  undissolved  matter  settle  and  note  red 
color  of  permanganic  acid.  This  is  a  good  test  for  Mn.  Pour  off  a 
little  of  the  clear  solution  into  water  to  see  color  better. 

(b)  Melt  in  an  iron  crucible  5  grams  solid  KOH  and  2.5  grams 
potassium  chlorate,  and  stir  in  gradually  2  grams  MnO2.  Heat  with 
stirring  till  the  mass  turns  solid  and  raise  the  temperature  with  full 
burner  flame,  and  continue  5  minutes.  When  the  mass  is  cold  dis- 
solve out  by  heating  with  the  crucible  nearly  full  of  water.  Pour  the 
solution  into  a  large  test  tube  and  let  settle.  To  one  portion  of  the 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  67 

green  solution  of  potassium  manganate,  K2MnO4,  add  dilute  sulfuric 
acid  till*  it  just  turns  red  forming  potassium  permanganate,  KMnO<. 
Try  changing  the  manganate  to  permanganate  in  another  portion  by 
pa-rsing  through  it  carbon  dioxide,  and  in  a  third  by  diluting  it  with 
much  water. 

147. — Dissolve  about  one-fourth  gram  of  oxalic  acid  in  water,  add 
5  c.c.  dil.  sulfuric  acid,  heat  to  about  80  degrees  and  add  a  little  at  a 
time  a  solution  of  potassium  permanganate  till  the  color  becomes  per- 
manent. Dissolve  about  one-half  of  a  gram  of  ammonium  ferrous  sul- 
fate  (143a),  add  sulfuric  acid  and  permanganate  as  above.  These  re- 
actions illustrate  oxidation  by  permanganic  acid  and  are  much  used 
in  quantitative  analysis.  Write  the  equations,  assuming  that  the  oxalic 
acid  is  oxidized  to  water  and  CO-  and  the  FeSO4  to  Fe-CSO^s. 
QUALITATIVE  SEPARATION  OF  THE  METALS 

The  following  scheme  of  analysis  is  prepared  for  first  year  stu- 
dents in  chemistry  to  be  used  in  the  separation  of  the  common  metals. 
Provision  is  not  made  for  every  contingency.  For  example,  it  is  as- 
sumed that  the  ions  to  form  insoluble  phosphates  in  Group  III  are 
not  present. 

148.— Group  I;  Ag,  Pb,  Hg:  Determine  with  test  paper  whether 
the  solution  is  neutral  or  only  slightly  acid.  If  neutral  add  one-tenth 
of  its  volume  of  HC1,  sp.  gr.  1.12.  If  strongly  acid,  neutralize  with 
ammonium  hydroxide  and  then  add  the  HC1.  The  purpose  is  to  have 
enough  acid  to  prevent  the  precipitation  of  BiOCl  in  this  group  and 
ZnS  in  the  next,  and  not  enough  to  prevent  the  precipitation  of  SnS 
and  CdS  in  group  II. 

If  no  precipitate  is  formed  pass  to  group  II.  If  one  is  formed  let 
remain  a  few  moments  and  filter.  Set  aside  the  filtrate  (1)  for  group 
(II).  Wash  twice  the  precipitated  chlorides  of  Ag,  Pb,  Hg2  with  small 
portions  of  cold  water.  Now  pour  through  the  filter  a  half  test  tube 
full  of  boiling  water.  Boil  and  pour  through  again.  Add  to  one-half 
of  filtrate  dil.  sulfuric  acid  and  let  stand.  A  white  precipitate  shows 
lead  sulfate.  To  the  other  add  potassium  chromate.  A  yellow  preci- 
pitate is  lead  chromate.  - 

Treat  the  remaining  precipitate  on  the  filter  with  about  5  c.c.  of 
NH\OH.  Pass  it  through  a  second  time.  Now  add  to  this  filtrate  an 
excess  of  dil.  nitric  acid,  making  sure  of  an  excess  by  mixing  and 
testing.  AgCl  is  obtained  if  silver  is  present.  The  blackening  of  the 
residue  left  on  filter  when  ammonia  was  added  shows  Hg2  present,  the 
black  substance  being  Hg  and  HgNH2Cl. 

Write  equations  for  all  reactions  that  occur  in  the  analysis  of 
Group  I. 

149.— Group  II;  As,  Sb,  Sn,  Hg,  Pb,  Bi,  Cu,  Cd:  Heat  filtrate  (1)  from 
group  (I)  nearly  to  boiling  and  pass  in  H2S  for  about  10  minutes  keep- 


68  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

ing  the  temperature  near  the  boiling  point  so  as  to  precipitate  arsenic 
from  arsenates  if  present.  Now  let  cool,  add  an  equal  volume  of  wa- 
ter and  pass  in  the  gas  for  several  minutes  to  insure  the  complete 
precipitation  of  cadmium  and  tin.  Even  then  it  is  well  to  let  the  pre- 
cipitate settle,  pour  off  a  few  drops  of  the  liquid,  add  water  and  more 
gas.  Disregard  a  white  precipitate  which  might  be  ZnS.  When  as- 
sured that  the  precipitation  is  complete  let  the  precipitate  settle  and 
pour  off  the  solution  through  a  filter.  At  once  boil  it  to  expel  H2S  and 
set  it  aside  as  filtrate  (2)  for  group  III. 

Wash  the  precipitate  four  times  by  decantation;  that  is,  pour  upon 
it  hot  water,  boil,  let  settle  completely  and  pour  off  the  wash  water. 
Drain  carefully  the  last  time  since  much  water  present  will  dilute  the 
nitric  acid  to  be  used. 

If  it  is  known  that  As,  Sb  and  Sn  are  absent  omit  the  bracketed  di- 
rections, and  boil  with  dil.  nitric  acid  as  directed  after  the  bracketed 
lines.  If  not  known  whether  they  are  present  or  not  proceed  as  di- 
rected within  the  brackets. 

[Now  heat  a  small  portion  of  the  washed  precipitate  with  about  3 
c.c.  of  yellow  ammonium  sulfide,  best  in  a  dish  on  a  water  bath  with 
stirring.  In  any  case  do  not  boil.  Filter  off  the  solution  and  add  to  fil- 
trate dil.  HC1  till  acid  throughout,  and  boil.  If  only  white  finely  di- 
vided sulfur  is  precipitated  As,  Sb  and  Sn  are  absent.  In  this  case  pro- 
ceed to  boil  the  remainder  of  the  precipitate  with  nitric  acid  as  after 
the  brackets.  If,  however,  a  flocculent  yellowish  precipitate  is  ob- 
tained one  or  more  of  the  above  metals  are  present.  Treat  the  re- 
mainder of  the  washed  precipitate  with  ammonium  sulfide  as  directed 
for  the  small  portion,  filter  and  set  aside  to  be  examined  for  As,  Sb 
and  Sn,  as  described  within  the  brackets  below. 

The  residue  consisting  of  sulfides  undissolved  by  ammonium  sul- 
fide must  be  well  washed  on  the  filter.  Punch  through  the  tip  of  the 
filter,  wash  the  sulfides  into  a  test  tube  and  let  settle.]  Pour  off  as 
much  as  possible  of  the  water,  add  about  10  c.c.  dil.  nitric  acid  and 
boil  persistently.  All  the  sulfides  except  that  of  Hg  are  dissolved.  A 
floating  residue  of  S  often  remains  and  is  to  be  disregarded.  A  heavy 
residue  which  settles  at  once  to  the  bottom  must  be  tested  for  Hg.  Fil- 
ter off  the  HgS,  add  to  the  nitrate  10  cc.  dil.  sulfuric  acid  and  begin  its 
evaporation  in  a  porcelain  dish.  To  confirm  the  presence  of  Hg  dis- 
solve the  sulfide  on  the  filter  by  pouring  upon  it  about  3  c.c.  of  hot 
con.  HC1  to  which  a  little  potassium  chlorate  has  been  added.  Heat  and 
pass  through  again  if  necessary.  Add  to  the  solution  bromine  water 
till  the  color  of  Br  persists,  boil  out  the  excess  and  add  a  few  drops  of 
a  solution  of  stannous  chloride.  A  white  precipitate  turning  to  gray 
and  perhaps  to  black  shows  Hg  present. 

Evaporate  the  nitric  acid  solution  containing  the  other  metals  till 


LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY  69 

heavy  white  fumes  of  sulfuric  acid  freely  come  off.  The  evaporation 
must  be  carried  far  enough  to  expel  all  the  nitric  acid.  (See  131)  Let 
the  dish  and  contents  cool,  then  add  10  c.c.  of  water  and  let  stand  5 
minutes.  A  white  precipitate  shows  lead  sulfate.  Filter  it  off  and  add 
to  filtrate  an  excess  of  ammonium  hydroxide,  making  sure  of  an  ex- 
cess by  shaking  to  mix  and  then  testing  with  turmeric  paper.  A  white 
precipitate  shows  Bi  but  its  presence  should  be  confirmed  thus:  Fil- 
ter it  off.  Dissolve  on  the  filter  in  a  few  drops  of  con.  HC1,  evaporate 
off  nearly  all  of  the  acid.  (Why?)  Pour  into  much  water.  A  white 
cloud  is  BiOCl. 

The  filtrate  from  the  bismuth  hydroxide  is  blue  if  copper  is  pres- 
ent. If  copper  is  absent  add  hydrogen  sulfide  to  the  filtrate  which  will 
give  yellow  CdS  if  Cd  is  present.  If  the  blue  color  shows  Cu  present 
test  for  Cd  by  one  of  the  following: 

(a)  Add  to  the  blue  solution  a  solution  of  potassium  cyanide  (very 
poisonous)  till  the  blue  color  disappears,  then  hydrogen  sulfide  which 
will  give  yellow  CdS  if  Cd  is  present. 

(b)  Add  to  the  blue  solution  dil.  sulfuric  acid  till  colorless  then 
card  teeth  or  steel  wool  and  boil  for  some  time.    The  iron  removes  the 
Cu.    Why?    Filter,  be  sure  that  the  filtrate  is  still  acid.    If  not  add  a 
slight  excess  of  sulfuric  acid  and  then  H2S  which  will  give  yellow  CdS 
if  Cd  is  present. 

[To  the  (NH4)2S  solution  containing  As,  Sb  and  Sn,  add  dilute  HC1  in 
excess,  which  re-precipitates  the  sulphides  if  present.  A  precipitate 
is  always  produced  owing  to  the  separation  of  sulphur.  If  the  sul- 
phides are  present,  however,  the  precipitate  is  more  highly  colored  and 
somewhat  flocculent.  Filter  and  wash  the  sulphides,  carefully  remove 
to  a  test  tube  and  heat  with  strong  HC1.  Sb2S3  and  SnS2  are  dissolved, 
while  AsaSs  is  not.  Dissolve  the  As2S3  in  hot,  strong  HC1  and  KCIO*, 
boiling  if  necessary,  best  in  a  small  flask  placed  in  the  evaporating 
closet.  Add  to  the  solution  NEUOH  in  excess  then  NH4C1,  filter  if  not 
clear  and  add  MgSO4,  which  will  give  on  standing  a  precipitate  of 
NH4MgAsO4  if  As  is  present. 

Boil  persistently  the  filtrate  from  the  arsenic  sulfide  with  card 
teeth.  The  Sb  is  deposited  on  the  card  teeth  from  which  it  can  be  de- 
tached in  scales  by  stirring  with  a  glass  rod,  and  the  tin  remains  in  so- 
lution as  SnCl2.  Filter  off  the  antimony  and  test  the  filtrate  for  tin 
with  HgCl2.  Wash  the  antimony  thoroughly,  dissolve  on  the  filter  with 
a  very  little  aqua  regia,  dilute  with  10  to  20  times  its  volume  of  water 
and  add  H2S  which  will  give  orange  yellow  Sb2S3L 

Write  equations  for  all  reactions  involved  in  Group  II  assuming 
that  all  metals  are  present  in  the  beginning  as  chloride. 

150— Group  III:  ..Fe,  Al,  Cr,  Co,  Hi,  Mn,  Zn:  To  a  small  portion 
of  filtrate  2  from  Group  II  add  ammonia  in  excess  as  proved  by  shak- 


70  LABORATORY  MANUAL  OF  GENERAL  CHEMISTRY 

ing  and  testing.  If  a  precipitate  is  formed  one  or  more  of  the  metals, 
Fe,  Al,  Cr  are  present.  If  no  precipitate  is  formed  add  H2S;  or,  if  a 
precipitate  was  formed  by  ammonia  filter  it  off  and  add  H2S  to  the  fil- 
trate. If  H2S  causes  a  precipitate  which  is  not  black  Co,  and  Ni  are 
absent.  A  black  precipitate  means  either  Co  or  Ni  or  both  and  they 
may  mask  Mn  and  Zn.  By  making  these  preliminary  tests  and  atten- 
tion to  the  following  much  work  and  time  may  be  saved. 

If  no  precipitate  was  formed  by  either  ammonia  or  H2S  proceed 
with  the  rest  of  the  filtrate  to  Group  IV.  If  a  precipitate  was  formed 
only  with  ammonia  there  is  no  need  to  test  for  Co,  Ni,  Mn,  Zn.  In  this 
case  add  ammonia  to  all  the  filtrate,  filter  off  the  hydroxides  of  Fe, 
Al,  Cr,  and  save  the  filtrate  for  Group  IV,  omitting  the  addition  of  H2S 
below  and  also  that  part  of  the  directions  applying  to  Co,  Ni,  Mn,  Zn. 
If  H2S  made  a  precipitate  while  ammonia  did  not  Al  and  Cr  are  ab- 
sent but  Fe  may  be  present  if  the  precipitate  was  black. 

The  following  scheme  provides  of  course  for  the  presence  of  all 
these  metals: 

To  filtrate  (2)  from  Group  II  add  ammonia  in  slight  excess  test- 
ing after  shaking,  then  add  about  2  c.c.  more.  Heat  the  solution  in  a 
flask  or  large  test  tube  nearly  to  boiling  and  pass  in  H2S.  Shake  and 
heat  frequently  to  make  the  precipitate  granular  and  more  easily  filt- 
ered and  washed.  When  the  precipitation  is  apparently  completed  fil- 
ter a  small  portion  and  pass  into  it  more  gas.  If  more  precipitate  forms 
add  more  H2S  to  the  whole  and  thus  proceed  till  complete  precipita- 
tion is  attained. 

Filter  through  a  fluted  filter  and  wash  with  hot  water.  Boil  the 
filtrate  (3)  till  all  hydrogen  sulfide  is  expelled,  filter  it  and  set  aside 
for  Group  IV.  As  soon  as  the  precipitate  produced  by  H2S  is  suffi- 
ciently washed,  (about  two  funnels  full  of  hot  water),  punch  through 
the  tip  of  the  filter  and  wash  most  of  the  precipitate  into  a  small 
flask  with  a  fine  stream  of  water  from  the  wash  bottle.  At  once  add 
considerable  excess  of  dil.  HC1.  Only  NiS  and  CoS  remain  undis- 
solved.  Filter,  reserve  filtrate  for  B  and  test  for  Ni  and  Co  in  A. 

A.  Dissolve  the  NiS  and  CoS  in  the  filter  in  about  3  c.c.  of  aqua 
regia  and  evaporate  the  solution  nearly  to  dryness  in  a  dish.  Add  about 
4  c.c.  water  then  NaOH  drop  by  drop  with  shaking  till  a  permanent 
precipitate  is  obtained.  Avoid  a  large  excess  of  NaOH.  Divide  the  li- 
quid and  suspended  precipitate  into  two  portions,  and  test  for  Co  and 
Ni  as  follows: 

1.  To  one  portion  add  about  2  cc.  acetic  acid,  30  per  cent,  and  15 
c.c.  of  a  25  per  cent  solution  of  potassium  nitrite  and  let  the  solution 
stand  half  an  hour.  A  yellowish,  white  granular  precipitate  is 


LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY  71 

2.  To  the  other  half  add  a  solution  of  potassium  cyanide  (dan- 
gerous) drop  by  drop  till  the  precipitate  just  dissolves.  It  is  essen- 
tial to  avoid  a  large  excess.  Warm  the  solution,  add  about  5c.c.  NaOH 
and  then  bromine  water  till  the  color  of  Br  persists.  A  black  precipi- 
tate often  forming  after  some  time  is  Ni(OH)8. 

B.  Evaporate  the  filtrate  from  the  CoS  and  MS  nearly  to  dryness, 
add  about  5  c.c.  water,  make  strongly  alkaline  with  NaOH  and  add 
gradually  with  stirring  about  2  grams  of  sodium  dioxide.    Boil  a  few 
minutes,  add  water  and  filter  off  the  Fe  and  Mn  hydroxides.    Save  the 
nitrate   for   C   and  wash   the   precipitate.     To   test  for   Mn  make   a 
bead  of  sodium  carbonate,  while  hot  touch  to  powdered  KC1O3.     Let 
cool,  take  up  with  it  a  little  of  the  precipitate  and  fuse  with  the  blow- 
pipe.   A  green  bead  shows  Mn  present.     To  confirm  the  presence  of 
Mn  and  test  for  Fe  remove  a  small  portion  of  the  precipitate  from  the 
filter,  dissolve  it  in  dilute  nitric  acid  and  add  a  solution  of  potassium 
ferrocyanide  which  will  give  a  dark  blue  color  if  Fe  is  present.    Now 
pour  upon  the  remaining  precipitate  about  3  c.c.  of  hydrogen  dioxide 
and  then  add  10  c.c.  nitric  acid,  sp.  gr.  1.2,  which  means  practically  1 
vol.  of  pure  concentrated  acid  diluted  with  2  parts  of  water.    Warm 
and  pass  through  again  if  necessary.    To  this  solution  add  a  little  at  a 
time  2-4  grams  of  lead  dioxide  or  red  lead  and  heat  to  gentle  boiling 
for  a  few  moments.    Let  the  suspended  matter  settle  when  the  liquid 
above  it  will  be  colored  red  by  permanganic  acid  if  Mn  is  present. 

C.  To  the  nitrate  from  the  Fe  and  Mn  hydroxides  which  may  con- 
tain Al,  Cr  and  Zn,  add  con.  nitric  acid  in  slight  excess  testing  with 
litmus  paper.    Add  5  c.c.  ammonium  chloride,  heat  and  add  ammonia 
in  slight  excess  as  shown  by  test  after  stirring.    A  white  flocculent 
precipitate  shows  aluminium  hydroxide.    Filter  it  off  and  acidify  the 
filtrate  with  acetic  acid,  mixing  and  testing  with  litmus  paper.     The 
liquid  is  yellow  if  chromate  radical  is  present.    Add  barium  chloride 
which  will  give  yellow  BaCrO4.    The  BaCl2  must  be  added  in  excess  to 
remove  all  the  CrO4  which  would  interfere  with  the  test  for  Zn.    Filter 
and  pass  through  several  times  if  necessary.    The  clear  filtrate  should 
not  be  yellow.    If  it  is  add  more  BaCl2  and  filter  again.    To  the  clear, 
colorless  filtrate  add  H2S  which  will  give  an  evident,  white  precipi- 
tate if  Zn  is  present.    Hydrogen  sulfide  always  produces  an  opales- 
cence  at  this  point  and  this  is  to  be  disregarded. 

GROUP  IT. 

151- — If  the  analysis  is  being  carried  through  all  the  groups  the 
solution  used  will  be  nitrate  (3),  in  which  case  it  should  be  made 
slightly  acid  with  dilute  HC1,  evaporated  to  two-thirds  of  its  volume, 
and  ammonium  hydroxide  added  till  alkaline  then  ammonium  carbon- 
ate. If  the  solution  is  an  original  one,  add  ammonium  chloride,  am- 


72  LABORATORY  MANUAL  OP  GENERAL  CHEMISTRY 

monium  hydroxide  till  alkaline  then  ammonium  carbonate  till  the  pre- 
cipitation is  complete.  To  determine  this  heat  the  solution  nearly  to 
boiling,  let  settle  and  add  a  few  drops  of  the  carbonate  to  the  clear  so- 
lution. Filter  and  to  a  small  portion  of  the  filtrate  add  sodium  phos- 
phate which  will  give  (NH4)MgPO4  if  magnesium  is  present.  Reserve 
the  remainder  of  the  filtrate  (4)  for  Group  V.  Wash  the  carbonates 
on  the  filter  and  dissolve  by  pouring  upon  the  filter  not  more  than  5 
c.c.  of  acetic  acid,  letting  it  run  through  into  a  clean  tube.  Run  through 
several  times  if  necessary  to  dissolve  all  the  carbonates.  To  a  small 
portion  of  the  acetic  acid  solution  add  twice  its  volume  of  a  saturated 
solution  of  calcium  sulfate.  Note  carefully  whether  the  precipitate  is 
formed  at  once  or  only  after  a  few  seconds.  Proceed  according  to  one 
of  the  following  as  required,  using  the  rest  of  the  acetic  acid  solution: 

(a)  If  no  precipitate  was  formed,  to  the  remainder  of  the  acetic 
acid  solution  add  ammonia  in  excess  and  ammonium  oxalate  which  will 
precipitate  calcium  oxalate. 

(b)  If  the  precipitate  was  slowly  formed,  only  Sr  and  Ca  can  be 
present.    To  the  acetic  acid  solution  add  a  solution  of  ammonium  sul- 
fate and  let  stand.     Filter  off  the  strontium  sulfate,  and  to  the  fil- 
trate add  ammonium  hydroxide  till  alkaline  then  ammonium  oxalate 
which  will  precipitate  CaC2O4,  if  Ca  is  present. 

(c)  If  the  precipitate  with  CaSO4  was  immediate  Ba  is  present, 
and  the  other  two  may  be.    To  the  acetic  acid  solution  add  an  excess 
of  pure  potassium  chromate  solution.    Filter  off  the  barium  chromate. 
Make  the  filtrate  alkaline  with  ammonia  and  add  ammonium  carbon- 
ate, which  will  precipitate  Ca  and  Sr  carbonates  if  they  are  present. 
Filter  off  the  carbonates,  wash  till  most  of  the  excess  of  chromate  is 
removed,  and  dissolve  on  the  filter  with  acetic  acid.    Test  a  small  por- 
tion of  the  filtrate  with  CaS04.    If  no  precipitate  forms  on  long  stand- 
ing Sr  is  absent  and  Ca  should  be  tested  for  as  in  (a).    If  Sr  is  pres- 
ent remove  it  from  the  remainder  of  the  acetic  acid  solution  with  am- 
monium sulfate,  let  stand  5  minutes,  filter  and  test  filtrate  for  Ca  as 
in  (b). 

GROUP  V. 

152.— Evaporate  a  part  of  filtrate  (4)  to  dryness  in  a  porcelain 
crucible;  heat  till  all  ammonium  salts  are  expelled,  and  test  with  the 
spectroscope  as  directed  in  lecture  for  Na,  K  and  Li.  If  uncertain  as 
to  the  spectrum  lines  and  flame  colors,  compare  with  spectra  of  the 
known  substances. 

To  test  for  NH4,  place  in  a  porcelain  crucible  a  small  amount  of 
lime,  add  enough  of  the  original  solution  to  moisten  it,  cover  with  a 
watch  glass  with  a  strip  of  moist  turmeric  paper  on  the  under  side, 
and  warm  gently.  If  NH4  be  present,  NH8  will  be  evolved  and  can  be 
recognized  by  its  action  on  the  turmeric  paper  and  by  its  odor. 


APPENDIX 


153. — Atomic  Weights  of  the  Common  Elements: 


Aluminum 

Al 

27.1 

Lead 

Pb 

207.20 

Antimony 

Sb 

120.2 

Lithium 

Li 

6.94 

Arsenic 

As 

74.96 

Magnesium 

Mg 

24.32 

Barium 

Ba 

137.37 

Manganese 

Mn 

54.93 

Bismuth  

Bi 

208.0 

Mercury  

ffff 

200.6 

Boron 

B 

11.0 

Molybdenum 

Mo 

96.0 

Bromine  

...  Br 

79.92 

Mckel  

IVi 

58.68 

Cadmium..  . 

....Cd 

112.40 

Nitrogen  

IV 

14.01 

Calcium  

...Ca 

40.07 

Oxygen  

0 

16.00 

Carbon  

...C 

12.00 

Phosphorus  

....p 

31.04 

Chlorine  

....Cl 

35.46 

Platinum  

Pt 

195.2 

Chromium 

....     Cr 

52.0 

Potassium 

K 

39.10 

Cobalt  

...Co 

58.97 

Silicon.... 

Si 

28.3 

Copper 

Cu 

63.57 

Silver 

Ag 

107.88 

Fluorine..  . 

F 

19.0 

Sodium.. 

...Na 

23.00 

Gold  

Au 

197.2 

Strontium  

Sr 

87.62 

Hydrogen.. 

..H 

1.008 

Sulphur     

s 

32.06 

Iodine  .    . 

I 

126.92 

Tin             ..    .    . 

Sn 

118.7 

Iron.    . 

Fe 

55.84 

Zinc     

...Zn 

65.37 

154. — Vapor  Pressure  of  Water  in  Millimeters  of  Mercury: 


Degrees  C. 

Pressure 

Degrees  C. 

Pressure 

15 

12.7 

26 

25.0 

16 

13.6 

27 

26.5 

17 

14.5 

28 

28.1 

18 

15.4 

29 

29.8 

19 

16.4 

30 

31.6 

20 

17.4 

31 

33.4 

21 

18.5 

32 

35.4 

22 

20.0 

33 

37.4 

23 

20.9 

34 

39.6 

24 

22.2 

35 

41.9 

25 

23.5 

iTendrixon 


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